Solubility of CO2 in amide-based Brønsted acidic ionic liquids

Solubility of CO2 in amide-based Brønsted acidic ionic liquids

J. Chem. Thermodynamics 57 (2013) 355–359 Contents lists available at SciVerse ScienceDirect J. Chem. Thermodynamics journal homepage: www.elsevier...

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J. Chem. Thermodynamics 57 (2013) 355–359

Contents lists available at SciVerse ScienceDirect

J. Chem. Thermodynamics journal homepage: www.elsevier.com/locate/jct

Solubility of CO2 in amide-based Brønsted acidic ionic liquids Dongshun Deng ⇑, Yanhong Cui, Dong Chen, Ning Ai ⇑ Zhejiang Province Key Laboratory of Biofuel, College of Chemical Engineering and Material Science, Zhejiang University of Technology, Hangzhou 310014, China

a r t i c l e

i n f o

Article history: Received 4 May 2012 Received in revised form 6 September 2012 Accepted 14 September 2012 Available online 10 October 2012 Keywords: Brønsted acidic ionic liquid (BAIL) Amide-based Carbon dioxide Solubility

a b s t r a c t Several hydrophilic amide-based Brønsted acidic ionic liquids (BAILs) were prepared by simple acid-base neutralization reaction of N,N-dimethylformamide (DMF), N,N-dimethylacetamide (DMAc), and N-methylpyrrolidone (NMP) with trifluoroacetic acid (TFA) or tetrafluoroboric acid (FBA). The solubility data of CO2 in these BAILs were determined at T = (303.15, 313.15, and 323.15) K and subatmospheric pressure using isochoric saturation method. With the same cation, CO2 solubility in TFA-based BAILs was higher than that in FBA-based ones. From the variation of solubility, expressed as Henry’s law constants, with temperature, the standard Gibbs free energy, enthalpy, and entropy changes of CO2 solvation were calculated. The solubilities of CO2 in these BAILs were apparently increased with increasing the molar volume of BAIL except for [DMFH][TFA]. Ó 2012 Elsevier Ltd. All rights reserved.

1. Introduction Ionic liquids (ILs) are nonvolatile liquid solvents composed of organic cations with organic or inorganic anions over a wide range of temperature. The versatile choices of cation and anion make the ILs as the ‘‘designer’’ substances for special application in chemical engineering process (e.g., absorbence and separation) [1]. In general, they have good solvency power for both inorganic and organic materials. Recently, ILs have been suggested as alternative liquid absorbents for CO2 in natural gas sweetening or in greenhouse gas control processes [2,3]. The performance of ILs as CO2 separation medium depends largely on three intrinsic properties such as nonvolatility, thermal stability and tunable chemistry, which make them be superior to common organic solvents in reducing environmental pollution, energy consumption and working exposure hazards [4–6]. Although a serial data on the solubility of CO2 in ILs were reported in literature, most of the published data are concentrated on the common dialkylimidazolium-based ILs [7–10], which involve several steps in the preparing procedure. Brønsted acidic ionic liquids (BAILs) can be simply synthesized by straightforward acid-base neutralization reactions without complex work-up [11]. The research of BAILs mainly emphasized on the potential applications for acid-catalyzed reaction or reaction-separation system [12,13]. Recently, Palgunadi et al [14] first reported the solubilities of CO2 in various lactam-based or amide-based BAILs and claimed that CO2 solubilities in these BAILs were relatively higher than those in imidazolium-based ILs. For further study on the

⇑ Corresponding authors. Tel.: +86 571 88320892. E-mail addresses: [email protected] (D. Deng), [email protected] (N. Ai). 0021-9614/$ - see front matter Ó 2012 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.jct.2012.09.023

dissolution of CO2 in various BAILs, in present paper, we reported the solubilities of CO2 in five BAILs obtained from neutralization of amides (DMF, DMAc and NMP) with equimolar TFA or FBA at T = (303.15, 313.15, and 323.15) K and subatmospheric pressure. The related thermodynamic property changes accompanying the dissolution were also presented. 2. Experimental 2.1. Chemicals The specifications of chemicals used are listed in table 1. The chemicals were commercially available and used without further purification. 1H-NMR spectra of the BAILs were recorded with a 500 MHz Bruker spectrometer in DMSO-d6 and calibrated with the tetramethylsilane (TMS) as the internal reference. Densities of pure BAILs at atmospheric pressure were carefully determined at T = (303.15, 313.15, and 323.15) K using a 5.567 ± 0.004 cm3 pycnometer (previously calibrated using double distilled water at 303.15 K) immersed in an oil-bath. The weight of BAIL was measured using electronic balance (Mettler-Toledo AL204) with an uncertainty of ±2  104 g. The pressure was monitored using a pressure transmitter (Yeli Engineering Hangzhou, WMB2088, 0 kPa to 200 kPa, with an accuracy of 0.25% full scale). 2.2. Synthesis of ionic liquids The amide-based ionic liquids of general type [amide][TFA] or [amide][FBA] were synthesized by simple neutralization reaction in the following way. All the BAILs were obtained as apparent low-viscous liquids and were found to be hydrophilic and stable

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TABLE 1 Specifications of chemical reagents. Chemical reagent

Source

Mass fraction purity (as received)

N,N-Dimethylformamide (DMF) N,N-Dimethylacetamide (DMAc) N-Methylpyrrolidone (NMP) Trifluoroacetic acid (TFA) Tetrafluoroboric acid (FBA) Carbon dioxide (CO2)

Sinopharm chemical reagent Sinopharm chemical reagent Sinopharm chemical reagent Sinopharm chemical reagent Kunshan dongmei chemicals Jingong Special Gas Co., Ltd.

Co., Co., Co., Co., Co.,

Ltd. Ltd. Ltd. Ltd. Ltd.

>0.992 >0.993 >0.992 >0.99 0.40 >0.9995

N,N-dimethylformamide tetrafluoroborate [DMFH][FBA]. 1HNMR(DMSO-d6): d(ppm) 15.41(s, 1H), 8.16(s, 1H), 2.87(s, 3H), 2.68 (s, 3H). N-methylpyrrolidone tetrafluoroborate [NMP][FBA]. 1HNMR(DMSO-d6): d(ppm) 15.46 (s, 1H), 2.92 (t, 2H), 2.57(s, 3H), 1.56 (m, 4H). 2.3. Experimental apparatus

FIGURE 1. Schematic diagram of the CO2 solubility apparatus. 1. magnetic stirrer and warmer; 2. CO2 gas equilibrium cell (EC); 3. CO2 gas reservoir (GR); 4. CO2 gas cylinder; V1–V4, stainless vavle; T. temperature controller; P. pressure transmitter.

in water or air. Purities of the BAILs were no less than 99% as determined from the 1H NMR results. Water content for all BAILs was less than 600 ppm as measured using Karl Fischer titration. N-methylpyrrolidone trifluoroacetate [NMP][TFA]. 9.91 g of N-methyl-2-pyrrolidine (0.1 mol) was added to 250 cm3 roundbottomed flask fitted with a reflux condenser and calcium chloride drying tube. The flask was immersed in a recirculating heated oil-bath and magnetic stirred. Then, 11.41g trifluoroacetic acid (0.10 mol) was dropped slowly into the flask with the temperature less than 313 K. The reaction lasted for another 3 h at 313 K to ensure the reaction had proceeded to completion. Subsequently, the impurities was evaporated under a reduced pressure (<300 Pa) at 348 K for several hours until the weight of the residue remained constant. The yield of [NMP][TFA] was 98% (20.89 g). 1 H-NMR (DMSO-d6): d(ppm) 15.50 (s, 1H), 2.89 (t, 2H), 2.56 (s, 3H), 1.58 (m, 4H). The following trifluoroacetate or tetrafluoroborate were synthesized similar as [NMP][TFA]. N,N-dimethylformamide trifluoroacetate [DMFH][TFA]. 1HNMR (DMSO-d6): d(ppm) 15.39(s, 1H), 8.14(s, 1H), 2.87(s, 3H), 2.71 (s, 3H). N,N-dimethylacetamide trifluoroacetate [DMAcH][TFA]. 1HNMR (DMSO-d6): d(ppm) 15.37(s, 1H), 2.70(s, 3H), 2.82 (s, 3H), 1.98(s, 3H).

CO2 solubility measurement was performed on the basis of the isochoric saturation method [15]. The glass-made apparatus, containing the equilibrium cell (EC, 2) with a magnetic stirrer and a gas reservoir (GR, 3) were connected as shown in figure 1. During the experiment, the EC and GR were immersed in the temperatureconstant water-bath. 2.4. Experimental procedure The volume of the EC (232.80 cm3) was directly calibrated by filling it with double distilled water at room temperature using titration column, with the accuracy of 0.01 cm3. The volume of the GR and the other parts of the system (640.01 cm3) were measured by the following procedure: with V3 closed while V1, V2, V4 opened, the whole system was evacuated, and then V1 was closed. The GR was pressurized with dried N2 and allowed to temperature equilibrium. The volume of the GR and the other parts of the system was calculated from the pressure drop observed when the V1 was opened using the ideal gas equation. During the experiment, the temperature of EC and GR was maintained at a certain value with a thermostatic water bath with a precision of ±0.05 K. The solubility of CO2 in ILs was determined as following: Each pure and dried IL was placed in the EC and degassed by vacuum pumping at 338K while stirring. The amount of each IL was determined by deduction weighing method. At a specified water-bath temperature, closed V3, opened V1,V2 and V4, the whole system was evacuated to pressure p1, then closed V1 and V4, opened V3, GR was loaded with CO2 from gas cylinder until the pressure reaches a scheduled value close to atmospheric pressure, recorded the pressure p2. Closed V2 and V3, opened V1, CO2 was brought into EC where absorbent was magnetic stirred

TABLE 2 Densities (q) and molar volumes (Vm) of BAILs at T = (303.15, 313.15, and 323.15) K. T = 303.15 K

q/(g cm3) [DMFH][TFA] 1.2460 [DMFH][FBA] 1.3408 [NMPH][TFA] 1.2323 [NMPH][FBA] 1.3175 [DMAcH][TFA] 1.2279

T = 313.15 K

T = 323.15 K

Vm/(cm3 mol1)

q/(g cm3)

Vm/(cm3 mol1)

q/(g cm3)

Vm/(cm3 mol1)

150.2

1.2329

151.8

1.2199

153.4

120.0

1.3345

120.6

1.3282

121.2

173.0

1.2261

173.8

1.2199

174.7

141.9

1.3117

142.5

1.3058

143.2

163.8

1.2180

165.1

1.2082

166.5

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D. Deng et al. / J. Chem. Thermodynamics 57 (2013) 355–359 TABLE 3 Experimental CO2 equilibrium pressure (p) and mole fraction (x2) in BAILs at various temperatures. T = 313.15 K 3

p/kPa

10 x2

18.77 32.74 43.94 55.57 72.69 92.02

3.81 6.65 8.82 11.25 14.69 18.65

20.28 37.93 50.23 60.43 72.16 91.88

1.92 3.61 4.72 5.71 6.77 8.61

19.12 32.54 43.22 52.37 71.56 91.02

3.96 7.01 8.84 11.28 14.95 19.12

21.13 40.12 53.53 67.58 80.25 92.33

0.93 1.76 2.37 2.98 3.54 4.08

18.17 32.66 44.06 53.99 72.18 88.53

3.67 6.48 8.66 10.64 14.12 17.73

T = 323.15 K 3

p/kPa 10 x2 [NMPH][TFA] + CO2 (MW: 213.15) 19.24 3.45 34.99 6.22 46.79 8.30 59.03 10.36 75.18 13.28 94.28 16.56 [NMPH][FBA] + CO2 (MW: 186.96) 21.72 1.75 39.56 3.19 52.79 4.24 63.59 5.13 75.03 6.11 93.14 7.48 [DMFH][TFA] + CO2 (MW: 187.12) 21.53 3.86 34.67 6.35 46.01 8.39 56.24 10.32 73.98 13.67 92.75 16.88 [DMFH][FBA] + CO2 (MW: 160.93) 23.25 0.91 41.93 1.64 55.79 2.18 70.45 2.75 83.73 3.28 95.12 3.71 [DMAcH][TFA] + CO2 (MW: 201.14) 19.02 3.27 34.57 5.94 46.74 8.03 57.58 9.89 74.12 12.78 89.01 15.31

vigorously. Progress of absorption was indicated by the slowly decrease of the pressure. It was assumed that equilibrium was reached after the pressure of the system had been constant for 4 h. The final pressure was recorded as p3. Thus, the amount of CO2 absorbed could be calculated from the differential between the total CO2 introduced into the GR and residual gaseous CO2 in the whole system. Since the pressure was subatmospheric pressure and the temperature was above room temperature, the absorbed CO2 amount was calculated using the ideal gas equation. The measurement was carried out under several temperatures to reveal the influence of temperature on the solubility of CO2 in each experimental IL.

3. Results and discussion Table 2 listed the measured densities and the calculated molar volumes for all the BAILs at atmospheric pressure and different temperatures. The densities of all the BAILs were less affected by the variation of temperature, but more influenced by the kinds of anions. The TFA-based BAILs exhibited lower density than the corresponding FBA-based ones, suggesting that trifluoroacetate anion create larger free volume within the BAILs. The densities among the BAILs with the same anion showed little differences. The molar volumes of the BAILs were in order of [NMPH][TFA] > [DMAcH][TFA] > [DMFH][TFA] > [NMPH][FBA] > [DMFH][FBA]. The CO2 solubilities in five BAILs at subatmospheric pressure and T = (303.15, 313.15, and 323.15) K were measured and illustrated in table 3, including CO2 mole fraction in liquid phase and CO2 equilibrium pressure above the absorbents.

p/kPa

103x2

20.51 37.15 49.73 62.72 77.94 96.98

3.08 5.64 7.55 9.41 11.78 15.01

24.16 41.64 55.24 66.51 78.31 95.82

1.69 2.90 3.83 4.58 5.46 6.63

23.46 36.73 48.74 60.02 75.72 95.62

3.72 5.85 7.55 9.47 12.02 15.01

25.43 43.55 65.42 73.14 86.46 97.13

0.88 1.51 2.27 2.53 2.99 3.37

21.46 36.85 49.55 61.23 76.53 92.17

3.22 5.46 7.35 9.05 11.47 13.51

20 16 12

103x2

T = 303.15 K

8 4 0

20

30

40

50

60

70

80

90

pCO / kPa 2 FIGURE 2. CO2 solubility, expressed as CO2 mole fraction (x2), as a function of CO2 equilibrium pressure (pCO2 ) at T = 303.15 K in: j, [DMFH][FBA]; ., [NMPH][ FBA]; 4, [DMAcH][TFA]; h, [NMPH][TFA]; e, [DMFH][TFA].

The amount of CO2 absorbed in BAILs can be calculated by the following equation,

nCO2 ¼ ððp2  p1 ÞV GR  ðp3  p1 ÞðV GR þ V EC  V IL ÞÞ=RT

ð1Þ

where VGR and VEC are the volumes of CO2 reservoir and equilibrium cell, respectively. VIL represents the volume of the experimental IL. The calculation was based on the ideal gas law because the

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D. Deng et al. / J. Chem. Thermodynamics 57 (2013) 355–359

TABLE 4 The experimentally determined Henry’s law constants (H) of CO2 in BAILs at various temperatures. T/K

H/MPa [NMPH][TFA] + CO2 4.94 5.67 6.56 [NMPH][FBA] + CO2 10.63 12.40 14.42 [DMFH][TFA] + CO2 4.76 5.46 6.35 [DMFH][FBA] + CO2 22.68 25.59 28.87 [DMAcH][TFA] + CO2 5.05 5.81 6.76

303.15 313.15 323.15 303.15 313.15 323.15 303.15 313.15 323.15 303.15 313.15 323.15

TABLE 6 Calculated standard Gibbs free energy (Ddis G0 ), enthalpy (Ddis H0 ) and entropy (Ddis S0 ) of dissolution of CO2 in BAILs at 0.1 MPa and different temperatures.

ra

T/K

Ddis G0 /(kJ mol1)

0.02 0.05 0.08

298.15 303.15 313.15 323.15

9.52 9.83 10.51 11.24

298.15 303.15 313.15 323.15

11.37 11.76 12.55 13.36

298.15 303.15 313.15 323.15

9.42 9.73 10.41 11.15

298.15 303.15 313.15 323.15

13.30 13.67 14.44 15.22

298.15 303.15 313.15 323.15

9.56 9.88 10.58 11.32

0.06 0.05 0.08 0.09 0.05 0.05 0.06 0.04 0.04

303.15 0.06 313.15 0.07 323.15 0.05 qffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi P measd 3 exp l 2 measd exp l a r¼ ðH H Þ =N , H ¼ 10 p=x2 , H is slope of the isotherm created from linear fit of CO2 mole fraction versus equilibrium pressure.

TABLE 5 The values of coefficients B0, B1 and B2 for Eq. (2) along with average absolute deviation of the fit (AAD).

a

BAILs

B0

B1

B2

AADa

[NMPH][TFA] [NMPH][FBA] [DMFH][TFA] [DMFH][FBA] [DMAcH][TFA]

16.805 12.838 19.761 13.031 18.68

6.6017103 3.5256103 8.452103 3.504103 7.7205103

8.1531105 3.1783105 1.101106 3.6322105 9.4818105

0.15 0.36 1.01 1.61 0.28

2

exp l

calcd

exp l

AAD ¼ 10 jH H j=N, H is Henry’s law constants in Table 4, H calculated from Eq. (2) at corresponding temperature.

calcd

experiment was carried out at subatmospheric pressure and above the room temperature. As shown in figure 2, the isothermal CO2 solubility at 303.15 K increases with increasing CO2 pressure. All the data exhibit linearity with zero intercept, which suggests the validity of the Henry’s law [16]. Such physical solubility can also be observed in many non-task-specific ILs [17–20]. Henry’s law constant can be used to reflect the solubility of CO2 in BAILs directly. In our experiments, because of the low equilibrium pressure and nonvolatility of BAILs, the fugacity of CO2 was approximately equal to equilibrium pressure [21], so the Henry’s law constants were determined from the slope of the isotherm created from linear fit of CO2 mole fraction versus equilibrium pressure. Henry’s law constants and the standard errors of the isotherm slopes are listed in table 4. The CO2 solubility in these BAILs is mainly determined by the anion and little affected by the structural distinction of the cation. The solubilities in the TFA-based BAILs are far higher than that in FBA-based ones with the same cation. When combining the Henry’s law constant in table 4 with the molar volume in table 2, it is evident that the solubility of CO2 in ILs enhances with increasing molar volume of IL except for [DMFH][TFA]. Indeed, this relationship between solubility and molar volume was similar with that in the imidazolium-based ILs [20,22,23] and other amide-based BAILs [14]. The particularity of CO2 solubility in [DMFH][TFA] may arise from the complex interaction between the CO2 and formoxyl group, the further investigation is being continued.

[NMPH][TFA] + CO2 9.42 10.17 11.59 12.93 [NMPH][FBA] + CO2 11.57 11.88 12.44 12.96 [DMFH][TFA] + CO2 8.87 9.88 11.81 13.62 [DMFH][FBA] + CO2 8.88 9.21 9.85 10.44 [DMAcH][TFA] + CO2 9.30 10.21 11.93 13.55

Ddis S0 /(J mol1 K1) 63.51 65.96 70.59 74.80 77.01 77.98 79.79 81.43 61.32 64.69 70.95 76.63 74.37 75.48 77.54 79.42 63.26 66.27 71.86 76.95

The behavior of Henry’s law constants as a function of temperature was correlated using an empirical equation (2) as following [24],

ln

is

Ddis H0 /(kJ mol1)



HðTÞ

 ¼

5

10 Pa

n X Bi ðT=KÞi

ð2Þ

i¼0

the optimized coefficients, Bi, obtained using a linear regression of multiple-variables calculation, are listed in table 5, along with the average absolute deviation (AAD) for each BAIL. Thermodynamic properties of dissolution of CO2 in five BAILs related to the Henry’s law constant can be calculated as follows:

Ddis G0 ¼ RT ln

Ddis H0 ¼ R

0

Ddis S ¼

  HðT; pÞ p0

ð3Þ

  @ lnðHðT; pÞ=p0 Þ @ð1=TÞ p

Ddis H0  Ddis G0 T

ð4Þ

! ð5Þ

in which Ddis G0 , Ddis H0 and Ddis S0 are the standard Gibbs free energy, enthalpy and entropy changes of dissolution of CO2 in five BAILs at the standard states pressure of p0 = 0.1 MPa, respectively. The thermodynamic property changes calculated from Eqs. (3)–(5) by incorporation of Eq. (2) and coefficients in table 5 were presented in table 6. For CO2 solvation in BAIL under all the conditions, the negative value of Ddis H0 indicates the process is exothermic, which means that the dissolution of CO2 in BAILs is favorable enthalpically. The standard Gibbs free enthalpy changes of approximate 9 kJmol1 to 12 kJmol1 in this work are consistent with previously reported values (10 kJmol1 to 14 kJmol1 for imidazolium-based ILs) [25,26], indicating the weak interaction of BAILs with CO2. The Ddis S0 is largely related to the BAIL organization surrounding CO2 [27]. The more negative entropy indicates the higher ordering degree when CO2 dissolves in the BAIL. As result, the Ddis G0 shows positive value.

D. Deng et al. / J. Chem. Thermodynamics 57 (2013) 355–359

4. Conclusion A series of amide-based BAILs were synthesized and the solubility data of CO2 in these liquids were determined at T = (303.15, 313.15, and 323.15) K and subatmospheric pressure. With the same cation, CO2 solubility in TFA-based BAILs was higher than that in FBA-based ones. The solubility of CO2 in these BAILs exhibited molar volume-dependent behavior except for [DMFH][TFA]. The experimental results were reduced to Henry’s law constants as a function of temperature. The standard Gibbs free energy, enthalpy, and entropy changes of CO2 dissolution in BAILs solutions were calculated. Acknowledgments The authors are grateful for the financial support by the Natural Science Foundation of Zhejiang Province (No. Y4100699) and the Natural Science Foundation of China (No. 21006095). References [1] J.D. Holbrey, R.D. Rogers, second ed., in: P. Wasserscheid, T. Welton (Eds.), Ionic Liquids in Synthesis, Wiley-VCH, Weinheim, Germany, 2007. [2] J.L. Anthony, S.N.V.K. Aki, E.J. Maginn, J.F. Brennecke, Int. J. Environ. Technol. Manage 4 (2004) 105–115. [3] J.E. Bara, D.E. Camper, D.L. Gin, R.D. Noble, Acc. Chem. Res. 43 (2010) 152–159. [4] R. Idem, P. Tontiwachwuthikul, Ind. Eng. Chem. Res. 45 (2006) 2413. [5] N. McCann, M. Maeder, M. Attalla, Ind. Eng. Chem. Res. 47 (2008) 2002–2009. [6] K.E. Gutowski, E.J. Maginn, J. Am. Chem. Soc. 130 (2008) 14690–14704. [7] L.A. Blanchard, D. Hancu, E.J. Beckman, et al., Nature 399 (1999) 28–29. [8] J.L. Anderson, J.K. Dixom, J.F. Brennecke, Acc. Chem. Res. 40 (2007) 1208–1216.

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JCT 12-237