Experimental solubility and thermodynamic modeling of CO2 in four new imidazolium and pyridinium-based ionic liquids

Experimental solubility and thermodynamic modeling of CO2 in four new imidazolium and pyridinium-based ionic liquids

Fluid Phase Equilibria 419 (2016) 67e74 Contents lists available at ScienceDirect Fluid Phase Equilibria j o u r n a l h o m e p a g e : w w w . e l...

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Fluid Phase Equilibria 419 (2016) 67e74

Contents lists available at ScienceDirect

Fluid Phase Equilibria j o u r n a l h o m e p a g e : w w w . e l s e v i e r . c o m / l o c a t e / fl u i d

Experimental solubility and thermodynamic modeling of CO2 in four new imidazolium and pyridinium-based ionic liquids Mohamed Zoubeik, Mohanned Mohamedali, Amr Henni* Acid Gas Removal Laboratory, Clean Energy Technologies Research Institute (CETRi), University of Regina, Regina, SK, Canada

a r t i c l e i n f o

a b s t r a c t

Article history: Received 25 October 2015 Received in revised form 3 March 2016 Accepted 10 March 2016 Available online 15 March 2016

The solubility of carbon dioxide (CO2) in four ionic liquids, 1-ethyl-3-methylimidazolium L-(þ)-lactate ([EMIM][LACTATE]), 3-methyl-1-propylpyridinium bis[(trifluoromethylsulfonyl]imide ([PMPY] [TF2N]), 1-(4-sulfobutyl)-3-methylimidazolium bis(trifluoromethanesulfonyl)imide ([(CH2)4SO3HMIm][TF2N]), 1-(4-sulfobutyl)-3-methylimidazolium hydrogen sulfate ([(CH2)4SO3HMIm][HSO4]), has been experimentally studied using the gravimetric microbalance method. (Carbon dioxide þ IL) systems were studied at (313.15, 323.15 and 333.15)K over a pressure range of 100 mbare20000 mbar. Experimental densities, Henry's Law constants, entropies and enthalpies of absorption were also reported. The results obtained showed that CO2 solubility diminished in the following sequence: [PMPY][Tf2N]>[EMIM] [LACTATE] > [(CH2)4SO3HMIm][TF2N] > [(CH2)4SO3HMIm][HSO4]. It was found that [PMPY][Tf2N] shows comparable CO2 solubility with ionic liquids that are considered promising such as [HMIM] [Tf2N], which makes this ionic liquid an attractive solvent for gas separation processes. CO2 solubility in the ionic liquids was well correlated using PengeRobinson equation of state with a quadratic mixing rule and the nonrandom two-liquid (NRTL) model. © 2016 Elsevier B.V. All rights reserved.

Keywords: CO2 solubility CO2 capture Ionic liquids EoS NRTL

1. Introduction Anthropogenic emissions of carbon dioxide represent one of the biggest environmental challenges of our generation [1]. Emission of carbon dioxide (CO2) is considered the most significant contributor to climate change, thus finding ways to minimize its release into the environment is the aim of many current research projects. Presently, the use of chemical aqueous amine solutions for CO2 capture is considered the most advanced technology for the removal of CO2 from flue gases. These methods are considered expensive due to the high energy required in the regeneration step [2]. Ionic liquids (ILs), as nonvolatile solvents, have shown great potential to act as a capture medium of CO2, and research interest has grown exponentially in the last few years in order to find the most efficient and environmentally friendly ionic liquid [3]. Ionic liquids possess physical properties (vapor pressure, melting point, and solubility) that can be systematically designed by selecting the proper cation and/or anion to achieve a given goal, hence the name “designer solvent” [4]. The main advantage of ionic liquids over other solvents

* Corresponding author. E-mail address: [email protected] (A. Henni). http://dx.doi.org/10.1016/j.fluid.2016.03.009 0378-3812/© 2016 Elsevier B.V. All rights reserved.

used for CO2 separation is their low vapor pressure and high thermal stability with the added benefit of minimal tendency for corrosion. Ionic liquids are therefore considered as promising for the future for clean, environmental and cost-efficient gas separation [3,5]. One of the most commonly investigated ionic liquids are the imidazolium-based ionic liquids along with other sulfonium, ammonium, and phosphonium derivatives [6]. Several ILs can be prepared by changing cations or anions or via chemical modifications of both of anion or/and cation. To improve CO2 solubility in ILs, several studies were performed to optimize the structure of the ILs in order to maximize CO2 solubility. Changing the intrinsic properties of ILs can generate a wide range of effects on the performance of ionic liquids, for instance, increasing the chain length of the alkyl group on the cation side results in an increase in CO2 solubility [7,8]. Anthony and coworkers have studied the influence of the anion on the gas solubility, and concluded that the nature of the anion has the most significant effect on gas solubility, regardless of the cation used [9]. Martinez et al. found that attaching a fluoro group to the alkyl part of an imidazolium cation can improve CO2 solubility by about 20% [10]. To further enhance CO2 solubility in ILs, another class of solvents called task specific ionic liquids was synthesized by amine modification of ILs to provide the reaction

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aspect to conventional ILs, and achieve an uptake of up to 0.5 mol CO2 per mol IL, or more, at atmospheric conditions [11]. Various review papers discuss the solubility in ILs with various chemical structures and in mixed ILs systems [12,13]. The objective of this work is to investigate the solubility of CO2 in four new ionic liquids: [PMPY][TF2N], [EMIM][LACTATE], [(CH2) 4SO3HMIm][TF2N] and [(CH2)4SO3HMIm] [HSO4] using a gravimetric microbalance (IGA). The selection of the aforementioned ionic liquids in particular was based on several characteristics, for instance the anion bis(trifluoromethylsulfonyl) imide which was proven to have high affinity toward CO2 due to the high fluorination content, while [EMIM][LACTATE] was chosen to investigate the influence of having long alkyl organic chain on CO2 solubility. Also we aimed to investigate the influence of varying the cation part to compare imidazolium (EMIM) and pyridinium based (pmpy) cations with the same anion. Furthermore, the CO2 solubility results obtained from this work were compared to other well-known ionic liquids in the literature. Finally, Henry's law constants, entropy, and enthalpy for the CO2 and ionic liquids were also calculated. The PengeRobinson equation of state with the quadratic mixing rule and the nonrandom two-liquid (NRTL) model were used to correlate the experimental data. The presence of water in the IL samples can influence the values of the thermophysical and transport properties such as density and viscosity, however the measurements performed in this study are for the samples as received from the suppliers which have water contents measured using Karl Fisher titration. The effects of water on density and solubility data were not investigated. Both the density and the solubility values reported here are therefore considered to be for the ILs with the water contents provided in Table 1. However, for the thermodynamic modeling calculations, only the CO2-IL system was considered, neglecting the presence of low water content which can be considered a reasonable approximation. 2. Experimental work 2.1. Materials Ionic liquids used in this work were obtained from SigmaeAldrich, io-li-tec and Solvionic, and are listed in Table 1 with their acronyms and purities. Research grade carbon dioxide (CO2), was purchased from Praxair, with a purity of 99.99 wt.%.

bi-distilled water. The U-tube was carefully cleaned and dried for 30 min at 353.15 K before injecting the ionic liquids. Approximately 2 mL of a sample was slowly injected inside the glass U-tube inside the apparatus. When the desired temperature was reached, the density was measured. The average of at least three measurements is reported (Table 2).

2.3. Solubility measurement Solubility was measured using an Intelligent Gravimetric Analyzer (IGA 003) from Hiden Analytical (Fig. 1). The gravimetric microbalance contains a sample bucket where the liquid is placed inside a pressure-vessel that is able to operate up to 20 bar and 500  C. For each experiment, a small amount of ionic liquid samples in the range of 60e90 mg was loaded into the sample container. Once the sample is loaded, the chamber was sealed. After stability is attained, the temperature was set at the degassing temperature of 348 K using an external water jacket. The sample is then dehydrated and degassed by completely evacuating the reactor using a diaphragm pump until the pressure reaches 20 mbar, followed by a turbo pump (Pfeiffer) to achieve a vacuum of about 10 mbar. The degassing step was continued for about 10 h to remove all traces of water and other volatile contaminants until a stable weight was achieved for about one hour, at which point the final weight was recorded. Temperature was then set at the absorption temperature using a water bath (Polyscience) with accuracy of 0.1 K. Three temperature settings used for this experiment were (313.15, 323.15 and 333.15) K. The sample temperature was measured with a type K platinum thermocouple (±0.1 K). When the isotherm temperature was reached, a set of desired pressure values, as well as parameters related to mass relaxation behavior, are set through the IGASwin software. The absorption process is then initiated by allowing CO2 via a mass flow controller (MFC) to reach a pre-set pressure inside the microbalance chamber. Any real-time weight change upon absorption was automatically recorded. Pressure and temperature are kept constant until equilibrium is reached. Then, the pressure was raised to the second data point of the isotherm, and this process was repeated for all other pressure measurements. Sufficient time of about 4 h was given to reach equilibrium and allow for weight stability at each pressure.

2.2. Density measurement

3. Modeling

The densities of ionic liquids used in this research were measured at different temperatures using an Anton Paar DMA 4500 digital density meter. The device allows for precision within 0.00001 g$cm3 and the uncertainty of the measurements was estimated to be 0.00005 g$cm3. The apparatus consists of a glass U-tube with a PT100 platinum resistance thermometer with an uncertainty of 0.01 K. The density meter was calibrated with air and

3.1. Calculation of Henry's law constants Using the concept and the relationship of the fugacity of the gas, Henry's law constant can be calculated. The fugacity was estimated from the experimental solubility data using the equation of state (PR-EoS). The Henry's Law constant is defined as:

Table 1 Characteristics of ionic liquids used in this work. Ionic liquid

Shorthand name

Supplier

Purity (wt.%)

Water content (wt.%)

Molecular weight

1-Ethyl-3-methylimidazolium L-(þ)-lactate 3-Methyl-1-propylpyridinium bis[(trifluoromethyl)sulfonyl] imide 1-(4-Sulfobutyl)-3-methylimidazolium bis(trifluoromethanesulfonyl)imide 1-(4-Sulfobutyl)-3-methylimidazolium hydrogen sulfate

[EMIM] [LACTATE] [PMPY] [TF2N]

Sigma Aldrich io-li-tec: ionic liquid technologies Solvionic

95 99

1 0.01

200.23 416.40

98

1

499.43

Solvionic

98

1

316.35

[(CH2)4SO3HMIm] [TF2N] [(CH2)4SO3HMIm] [HSO4]

M. Zoubeik et al. / Fluid Phase Equilibria 419 (2016) 67e74

69

Table 2 Experimental densities of pure ionic liquids measured at 1.013 bar.

r (g$cm3)

T (K)

278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15 333.15 338.15 343.15 348.15 353.15

[PMPY ][Tf2N]

[EMIM][LACTATE]

[(CH2)4SO3HMIm][TF2N]

[(CH2)4SO3HMIm][HSO4]

1.46723 1.46243 1.45762 1.45281 1.44803 1.44328 1.43854 1.43381 1.42911 1.42443 1.41979 1.41516 1.40980 1.40596 1.40140 1.39685

1.15745 1.15347 1.14977 1.14611 1.14246 1.13889 1.13543 1.13195 1.12851 1.12507 1.12167 1.11851 1.11798 1.11166 1.10839 1.10514

e 1.59792 1.59296 1.58785 1.58258 1.57787 1.57318 1.56852 1.56387 1.55923 1.55460 1.55001 1.54556 1.54112 1.53668 1.53223

e 1.45030 1.44702 e 1.43708 1.43368 e 1.42676 e 1.42051 e 1.41428 e 1.40807 e 1.40180

3.2. Critical properties calculations To calculate the parameters for the modeling equations, the critical temperature (Tc), critical pressure (Pc), and the acentric factor (u) of both components, CO2 and ionic liquids, are needed. These values are obtained by estimation using the classical LyderseneJobackeReid modified method [14]:

Tb ðKÞ ¼ 198:2 þ Tc ðKÞ ¼

AþB

Pc ðbarÞ ¼

P

X

nDTb Tb

nDTC  ð

(3)

P

nDTC Þ2

M P C þ ð nDPC Þ2

  X nDVC VC cm3 $mol1 ¼ D þ

Fig. 1. Schematic of the gravimetric microbalance.

(4)

(5)

(6)

In the equations above, M is the molecular weight of the ionic liquids used. The critical property contributions can be found in the literature [15]. The acentric factor is calculated as follows:

Hi ðT; PÞ ¼ lim xi/0

fiL xi

u¼ (1)

fLi

where, is the fugacity of the gas dissolved in the liquid phase. Since the fugacity of the gas in the liquid phase must be equal to the fugacity of the gas in the gas phase and approximating the gas phase fugacity as the gas phase pressure, the following form of Henry's law can be obtained:

Pi ¼ Hi ðTÞxi

    ðTb  43ÞðTc  43Þ ðTc  43Þ Pc Pc log log  ðTc  Tb Þð0:7Tc  43Þ ðTc  Tb Þ Pb Pb   Pc þ log 1 Pb

(7)

The equation for the acentric factor needs the calculated critical properties and the calculated normal boiling temperature as input parameters. The normal boiling temperature is obtained at the normal boiling pressure of (Pb ¼ 1.013 bar).

(2)

where, Pi is the partial pressure of the gas and Hi(T) will have units of pressure and is inversely proportional to the mole fraction of gas in the liquid. Henry's constants are proportionality constants that are used to relate the partial pressure of a gas to the gas solubility in a liquid state at infinitely dilute conditions [14]. Henry's law constant was then found by fitting the data to a second order polynomial with R2 > 0.999. The limit of the mole fraction of CO2 as it approaches zero pressure was used to obtain the Henry's law constant for each temperature.

3.3. Peng Robinson EoS The experimental CO2 solubility was correlated using PengeRobinson equation of state (PR-EoS) [16]. The PR-EOS parameters are given by the following equations:



RT aðTÞ  V  b VðV þ bÞ þ bðV  bÞ

(8)

70

M. Zoubeik et al. / Fluid Phase Equilibria 419 (2016) 67e74 2.0 [PMPY][TF2N] [EMIM][LACTATE] [(CH2)4SO3HMIm][TF2N] [(CH2)4SO3HMIm][HSO4]

1.8

p (g/cm3)

1.6

 i2 h ai ðTÞ ¼ 1 þ mi 1  Tri0:5

(11)

mi ¼ 0:37464 þ 1:54226wi  0:26992w2i

(12)



1.4

1.2



XX XX

 0:5   xi xj ai aj 1  kij

xi xj

(13)

bi þ bj 1   2 1  Iij

(14)

1.0 260

280

300

T(K)

320

340

360

3.4. NRTL activity coefficient method Fig. 2. Experimental density of pure ionic liquids.

The general nonrandom two-liquid (NRTL) equation was also used to correlate the vaporeliquid equilibrium of IL containing systems. Using AspenPlus, NRTL generates the liquid activity coefficients. The activity coefficients, gi, were correlated by the following equations based on the NRTL model:

Table 3 Temperature-dependent density correlations for the studied ionic liquids. Ionic liquids

Density (g$cm3)

AAD (%)

[EMIM] [LACTATE] [PMPY ][Tf2N] [(CH2)4SO3HMIm][TF2N] [(CH2)4SO3HMIm][HSO4]

r (g$cm3) ¼ 1.1601  0.0007  T [ C] r (g$cm3) ¼ 1.1416  0.0009  T [ C] r (g$cm3) ¼ 1.6016  0.0009  T [ C] r (g$cm3) ¼ 1.4533  0.0006  T [ C]

0.02 0.06 0.09 0.14

" lng1 ¼

x22

lng2 ¼

x21



t21 "



t12

G21 x1 þ x2 G21 G12 x2 þ x1 G12

2 þ

(15)

ðx2 þ x1 G12 Þ2

2 þ

!#

t12 G12

!#

t21 G21

(16)

ðx1 þ x2 G21 Þ2

Table 4 Critical properties. Ionic liquids

MW (g$mol1)

Tb (K)

Tc (K)

Pc (bar)

Vc (cm3$mol1)

u

ZC

[PMPY][Tf2N] [EMIM][LACTATE] [(CH2)4SO3HMIm][TF2N]] [(CH2)4SO3HMIm][HSO4]

416.36 200.23 499.43 316.4

839.8 693.4 1097.6 1017.6

1234.2 912.7 1612.8 1433.0

27.55 28.24 32.7 25.88

964.7 620.1 1070.1 744.8

0.3070 0.9702 0.377 0.8437

0.2591 0.2260 0.2615 0.3602

ai ¼ ai 0:45724

bi ¼ 0:07780

where,

R2 Tci2

(9)

2 Pci

RTci Pci

(10)

G12 ¼ expðf12 t12 Þ G21 ¼ expðf12 t21 Þ t12 ¼ ðg12  g22 Þ=RT t21 ¼ ðg21  g11 Þ=RT

20

The parameters g12 and g21, are equivalent (g12 ¼ g21). t12, t21 and a12 are three binary parameters adjusted to the experimental solubility data of ionic liquids.

18

Pressure (bar)

16 14 12

4. Results and discussion

10 8

4.1. Density

6 4

[Emim][LACTATE] ( this work )

2

[Emim][Ac] (Shiflett et al.(2008))

0 0.1

0.2

0.3

0.4

Mole fraction of CO2 Fig. 3. Comparison of [Emim][Ac] and [Emim][LACTATE] at 50  C.

0.5

Fig. 2 shows the experimental measurements of the density for the four ionic liquids. Density values decreased linearly with increasing temperature, with [(CH2)4SO3HMIm] [TF2N] showing the highest density while the lowest density was recorded for [EMIM][LACTATE] to be around 1.1000 g/cm3. The temperature-dependent correlations of density for each ionic liquid are expressed as a linear function as shown in Table 3.

M. Zoubeik et al. / Fluid Phase Equilibria 419 (2016) 67e74 40000

Pressure(M bar)

4.2. Critical properties

[bmim][Ac] (Shiflett et al., 2008) [BMIM] [PF6] (Shiflett and Yokozeki, 2005) [(CH2)4SO3HMIm][TF2N] (This work) [(CH2)4SO3HMIm][HSO4] (This work) [PMPY ][TF2N] (This work) [EMIM][LACTATE] (This work)

35000 30000

71

The estimated critical properties, normal boiling temperatures, and the acentric factors of the four ionic liquids used in this work are shown in Table 4 below.

25000 20000

4.3. CO2 solubility

15000 10000 5000 0 0

10

20

30

40

50

Mole fraction of CO2 in ionc liquids (%) Fig. 4. Comparison of CO2 solubility of the two ionic liquids with the same anion at 323.15 K.

The average absolute deviations (AADs) between the experimental and calculated densities of [EMIM] [LACTATE], [PMPY][Tf2N], [(CH2)4SO3HMIm][TF2N] and [(CH2)4SO3HMIm][HSO4] are 0.02, 0.06, 0.09 and 0.14 respectively.

CO2 solubility was measured at (313.15, 323.15 and 333.15) K at different pressures up to 20 bar as reported in Table 5. However, for [(CH2)4SO3HMIm][TF2N] and [(CH2)4SO3HMIm] [HSO4] the solubility was only measured at 313.15 and 323.15 because the viscosity appeared high prior to CO2 absorption. The nature of some ionic liquids is for viscosity to increase further upon CO2 fixation which is attributed to the formation of hydrogen bonding and the dominant Coulombic force [18]. As a result, once CO2 solubility was measured at 313.15 and 323.15 K and found to be quite low, the experiment was terminated due to high viscosity that was evident. It was clear that these two ionic liquids would have some drawbacks as potential IL for industrial use due to their viscous nature. Solubility of CO2 in the ionic liquids decreased as temperature increased and pressure decreased as reported in Table 5. CO2

Table 5 Experimental solubility (P, x) data for [PMPY][TF2N] + CO2, [emim][LACTATE] + CO2, [(CH2)4SO3HMIm][TF2N]] + CO2, [(CH2)4SO3HMIm][HSO4] systems at 313.15, 323.15 and 333.15 K.a [PMPY][TF2N] + CO2 system

[EMIM][LACTATE] + CO2 system

[(CH2)4SO3HMIm][TF2N]] + CO2 system

T/K

P (mbar)

102 x(CO2)

T/K

P (mbar)

102 x(CO2)

T/K

P (mbar)

102 x(CO2)

313.15

99.0 500 1001 1999 3998 7000 8999 10,000 10,999 12,998

0.238 1.09 2.12 4.31 8.34 13.7 17.1 18.8 20.4 23.5

313.15

999 2001 3999 6999 8999 9998 11,000 12,997 14,999 17,000

8.7 11.6 15.2 18.9 21.0 21.8 22.7 24.2 25.6 27.1

313.15

98 500 999 8998 9998 11,000 12,998 14,999 16,999 19,000

0.1328 0.1319 0.734 13.5 14.6 15.9 18.3 20.8 23.0 25.4

14,998

26.3

18,998

28.3

323.15

16,998 18,997

29.0 31.7

323.15

333.15

a

96 500 998 1999

0.229 0.928 1.87 3.79

3998

7.25

7001 9000 9996 10,999 13,002 14,998 17,007

11.8 14.7 16.1 17.5 20.2 22.8 25.3

19,008

27.6

100 500 999 2000 3999 6997 9001 9999 10,999 12,997 15,000 16,995 18,996

0.2.19 0.744 1.54 3.20 6.06 10.0 12.6 13.7 14.9 17.2 19.3 21.4 23.2

323.15

999 3998 4998 6999 9998 18,998

333.15

7.14 13.0 14.3 16.4 19.1 25.0

500

4.30

998 2000 3996 7000 9001 9998

5.97 8.29 11.2 14.3 15.9 16.7

Standard uncertainties u are u(T) ¼ 0.1 K, u(P) ¼ 8 mbar, and u(xCO2) ¼ 0.006 mol%.

98

0.158

500 1000

0.621 1.22

3998 7000 8999 10,001

5.57 9.09 11.59 12.58

15,002

18.2569

[(CH2)4SO3HMIm][HSO4] + CO2 system 313.15 99 0.0764 498 0.1753 999 0.2961 6999 2.53 11,000 4.82 12,997 5.52 323.15

101

0.056

500 9001 9999 10998.7 12,998 15,000 17,000

0.116 3.19 3.65 4.08 4.73 5.48 6.14

72

M. Zoubeik et al. / Fluid Phase Equilibria 419 (2016) 67e74

Table 6 Henry's Law constants and enthalpies and entropies of absorption for CO2 in the studied ionic liquids. Ionic liquids

H (bar)

[PMPY][Tf2N] [EMIM ][LACTATE] [(CH2)4SO3HMIm][TF2N] [(CH2)4SO3HMIm] [HSO4]

313.15 K

323.15 K

333.15 K

43.7 46.2 58.8 274

52.1 54.4 70.9 301.4

60.1 64.8

Table 7 Binary interaction parameters of the standard PR-EoS for the ionic liquids þ CO2 system. IL þ CO2

T (K)

[PMPY][Tf2N]

Binary interaction parameter

313.15 323.15 333.15 313.15 323.15 313.15 323.15

[(CH2)4SO3HMIm][TF2N] [(CH2)4SO3HMIm] [HSO4]

AAD%

kij

Iij

0.062 0.044 0.052 0.040 0.043 0.160 0.140

0.027 0.465 0.316 0.422 0.370 0.814 0.710

3.15 6.70 5.45 7.36 5.91 4.50 5.40

Ds (J$mol$K1)

13.8 14.6 15.7 8

42.9 45.7 49.5 25

solubility decreased in the following order: [PMPY][Tf2N] > [EMIM] [LACTATE] > [(CH2)4SO3HMIm][TF2N] > [(CH2)4SO3HMIm][HSO]. All ionic liquids had an acceptable percentage of impurity loss (less than 5 wt. %) except [(CH2)4SO3HMIm] [TF2N] with 23.10 wt. %, and [(CH2)4SO3HMIm][HSO4] with 8.60 wt. %. These two hydrophobic ionic liquids had more than the expected weight loss which might be due to their absorption of water when exposed to ambient atmosphere, and some acid evaporation loss (reversed reaction of formation) due to the vacuum initiated and high temperature. However, this would not affect the margin of error in terms of CO2 solubility assessment as the IGA does use weight change to assess the amount of CO2 absorption, however this occurs once the weight is stabilized and all impurities are removed, before the introduction of CO2. 4.4. The effect of changing the cation

2000 1800 1600 Pressure (kPa)

Dh (kJ$mol1)

1400 1200 1000 800 600 400 200 0 0

0.05

0.1

0.15 0.2 0.25 CO2 Solubility (xCO2)

0.3

0.35

Fig. 5. Comparison of isothermal solubility data of CO2 in different ionic liquids. The markers are the experimental data at each temperature as following C 313.15 K, 323.15 K, and :333.15 K. The lines represents the PR model estimations where e (Blue) is for [PMPY] [TF2N], (Red) for [HMIM] [TF2N], and (Green) is for [HMIM] [HSO4]. (For interpretation of the references to color in this figure legend, the reader is referred to the web version of this article).

When comparing the two ionic liquids with the same anion ([Tf2N]), changing the cation can obviously change CO2 solubility. At 313.15 K and at 19 bar [PMPY][TF2N] has a solubility of 0.3166 (mole fraction) and [(CH2)4SO3HMIm][TF2N] has a solubility of 0.2537 (mole fraction). [PMPY] is a pyridinium based cation and [(CH2)4SO3HMIm] is an imidazolium based cation which might be responsible for their difference in CO2 solubility. However, this finding is inconsistent with trends in the literature, for example, it is expected that [(CH2)4SO3HMIm] would provide a higher CO2 solubility due to its long alkyl chain length and its higher density and thus free volume, however this does not hold true for this case [7]. Therefore, there might be other factors that affect CO2 solubility that are yet to be discovered. 4.5. The effect of changing the anion When comparing the two ionic liquids with the same cation [(CH2)4SO3HMIm], changing the anion has a noticeable effect on CO2 solubility. At 313.15 K and 13 bar, [(CH2)4SO3HMIm] [TF2N]

(Pexp - Pcal)/Pexp *100%

15

2000

Pcalc (kPa)

1500 1000 500 0 0

500

1000

1500

2000

Pexp (kPa) Fig. 6. Parity Plot of the experimental versus calculated pressures for all the ionic liquids and all temperatures.

10 5 0 0

0.05

0.1

0.15

0.2

0.25

0.3

0.35

-5 -10 -15 CO2 Solubility (xCO2)

Fig. 7. Deviation of pressure at different solubility values for all the ionic liquids and all the temperatures.

M. Zoubeik et al. / Fluid Phase Equilibria 419 (2016) 67e74

73

Table 8 Binary interaction parameters of the NRTL model for the ionic liquids þ CO2 system. Ionic liquids

Binary interaction parameters

[PMPY][TF2N]

(g12 (g21 (g12 (g21 (g12 (g21

[(CH2)4SO3HMIm] [TF2N] [(CH2)4SO3HMIm] [HSO4]

     

g22)/R g11)/R g22)/R g11)/R g22)/R g11)/R

T (K)

Linear function

313.150

323.150

333.150

186.869 23.244 1435.801 496.940 32.830 653.800

152.945 73.629 523.701 629.901 72.030 246.701

119.020 124.013

has a solubility of 0.1829 (mole fraction) while [(CH2)4SO3HMIm] [HSO4] has a solubility of 0.0519 (mole fraction). Clearly, the fluorinated anion dramatically increases CO2 solubility. Cadena et al. showed that the nature of anion has the greatest influence on the solubility of CO2 and the bis(trifluoromethylsulfonyl)-imide anion [Tf2N] had the greatest affinity for CO2 [17]. This increase in CO2 solubility is in part explained by the strong coulombic interactions responsible for the organization of the liquid [18]. In addition, high CO2 solubility, as seen especially with the [TF2N] anion, is attributable to the fluoroalkyl groups in [Tf2N] which are known to be CO2-philic [19]. This may be the result of the favorable interactions between the negative fluorine atoms of these anions and the positive charge on the carbon in carbon dioxide [19,20]. 4.6. Comparison of all ionic liquids studied in this work

Pressure (bar)

CO2 solubility for four different ionic liquids was studied in this work. CO2 solubility decreased in the following order [PMPY] [Tf2N] > [EMIM] [LACTATE] > [(CH2)4SO3HMIm] [TF2N] > [(CH2) 4SO3HMIm][HSO4]. As shown in Fig. 4, the ionic liquid, [EMIM] [LACTATE], has a different curve shape than the rest of the ionic liquids. One notable fact is that it has a high CO2 solubility at low pressures. This is similar to Shiflett's ionic liquid, [EMIM][Ac], as seen in Fig. 3 [21]. It was demonstrated that it had high solubility even at very low pressure, which means that [EMIM] [LACTATE] acts as a physical and chemical solvent when interacting with CO2. As shown by Blath et al., IL with carboxylic anions such as [EMIM] [pivalate], [EMIM][OAc] and [EMIM] [benzoate] show chemoabsorption tendencies [22]. In addition, [LACTATE] contains a carbonyl group which is known to increase CO2 solubility by

(g12 (g21 (g12 (g21 (g12 (g21

     

g22)/R g11)/R g22)/R g11)/R g22)/R g11)/R

¼ ¼ ¼ ¼ ¼ ¼

AAD%

[3.3424  T(K)]  1299.221 [5.03844  T(K)] þ 1554.54 [91.2  T(K)] þ 3000 [13.3  T(K)] þ 3690 [68.66  T(K)]  21534.8 [31.88  T(K)] þ 10056.81

0.9 1.1 1.1

providing free electrons on the oxygen and allowing them to interact with the Lewis acidic carbon of CO2 [7]. 4.7. Comparison of this work with ionic liquids from the literature The ionic liquid [PMPY][Tf2N] shows promising performance when it comes to CO2 absorption as it has a similar solubility pattern to some ionic liquids published in the literature such as [HMIM][Tf2N] [23] which is attributed to the high fluorination content. In addition, we can also compare fluorinated ionic liquid studied in this work with [BMIM][PF6] which has a highly fluorinated anion; however, it has a lower CO2 solubility in comparison to the [Tf2N] anion [24]. Solubilities of the ionic liquids studied in this work are compared to [BMIM] [Ac] which clearly has a much higher CO2 solubility as shown in Fig. 4. [BMIM] [Ac] is known in the literature for having one of the highest CO2 absorption to date [21]. The high solubility results from the chemo-absorption properties of the ionic liquid which is unlike any other ionic liquid as shown in Fig. 4 [22]. [BMIM] [Ac] was found to form a carboxylate intermediary compound via a carboxylation reaction that was irreversible even at 363 K [22]. The chemical and physical absorptive capacity of [BMIM] [Ac] is shown by the high solubility at low pressures. This type of reaction and irreversible nature of [BMIM] [Ac] þ CO2 system and loss of solvent is a concern in industrial application despite its high CO2 absorptive capacity. 4.8. Henry's law constants Henry's Law constants for CO2 in the ionic liquids used in this work at the three different temperatures are presented in Table 6. Henry's Law constants were found by fitting the data to a first and second order polynomial. The ionic liquid with the lowest Henry's law constant is [PMPY][Tf2N] which was already been shown to be the ionic liquid with the highest CO2 solubility in the Results Section (Table 5). 4.9. Enthalpy and entropy calculations Enthalpy and entropy values for CO2, in the studied ionic liquids, are reported in Table 4. In the table, negative enthalpy values show that CO2 has a strong interaction with most of the ionic liquids perhaps the strongest interaction is seen with [(CH2)4SO3HMIm] [TF2N] [24]. The negative values for entropy indicate a higher degree of ordering as CO2 dissolved in these ionic liquids [24].

0

0.05

0.1

0.15

0.2

0.25

0.3

0.35

CO2 Solubility (xCO2) Fig. 8. Comparison of isothermal solubility data of CO2 in different ionic liquids. The markers are the experimental data at each temperature as following C 313.15 K, 323.15 K, and :333.15 K. The lines represents the NRTL model estimations e (Blue) for [PMPY] [TF2N], (Red) for [HMIM] [TF2N], and (Green)[HMIM] [HSO4]. (For interpretation of the references to color in this figure legend, the reader is referred to the web version of this article).

4.10. Thermodynamic modeling of CO2 solubility PR-EOS was used to correlate the experimental solubility data of the ionic liquids. Binary interaction parameters were evaluated by regression of the experimental data with PR-EOS and the quadratic mixing rule at different temperatures using a code developed with MATHEMATICA® software. The average absolute deviations in

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M. Zoubeik et al. / Fluid Phase Equilibria 419 (2016) 67e74

percentage (AAD%) between the calculated and the experimental equilibrium pressures which are defined below, were calculated for each ionic liquid and temperature.

AAD% ¼

exp 100 X Pi  Picalc N Piexp

of a Graduate Research Fellowship (GRF) for the second co-author. In addition, the first co-author would like to thank the Canadian Bureau of International Education (CBIE) and the Libyan government for a graduate scholarship.

(17) References

where, N is the number of equilibrium data points at each temperature, Pexp and Pcalc are the experimental equilibrium pressure and the calculated pressure, respectively. Table 7 summarizes the values for the interaction parameters along with the AAD% at each temperature. The predictions of PR-EOS for the ionic liquids at different temperatures is shown against the experimental data in Fig. 5. And as can be observed from the AAD% in Table 7, the PR-EOS satisfactorily correlated the experimental data for the temperature range studied. Fig. 6 shows the parity plot of the calculated versus the experimental equilibrium pressures for all the ionic liquids and temperatures studied. The deviations in the calculated pressure are shown versus the CO2 solubility in Fig. 7. The deviation was kept in a range of 10%. ASPEN plus® was used to perform the regression of the experimental data with NRTL model to evaluate the model parameters as shown in Table 8, with maximum AAD% of about 1.1, while the NRTL estimations versus experimental values is shown in Fig. 8 for all the ILs studied in this work at different temperatures. Thermodynamic modeling results for [EMIM][LACTATE] were not reported as the IL appeared to interact chemically with CO2 during the absorption hence the PR-EoS inherently cannot describe the chemical interaction behavior between the gas molecules and the liquid absorbent. The values obtained from our code for kij and Iij were in the order of magnitude of 8, therefore, these values were not reported in Table 7. 5. Conclusion CO2 solubility for four different ionic liquids was studied in this work. The CO2 solubility was measured at several fixed temperatures (313.15, 323.15, 333.15 K) and at pressures up to 20 bar. CO2 solubility decreased in the following order [PMPY][Tf2N] > [EMIM] [LACTATE] > [(CH2)4SO3HMIm][TF2N] > [(CH2)4SO3HMIm] [HSO4]. The most promising ionic liquid in this study is [PMPY] [Tf2N], which showed high CO2 solubility similar to other wellknown ionic liquids published in the literature such as [HMIM] [Tf2N]. [EMIM][LACTATE] showed very high solubility at low pressure, but could not be correlated with the PR equation of state, as it behaved like [BMIM][Ac], known to react chemically with CO2. Acknowledgment The authors acknowledge the support of the Acid Gas Removal Laboratory (AGRL) at the University of Regina as well as the Faculty of Graduate Studies and Research for financial support in the form

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