Article pubs.acs.org/jced
CO2 Solubility in Biodegradable Hydroxylammonium-Based Ionic Liquids Stepan Bazhenov,† Mahinder Ramdin,‡ Alexey Volkov,† Vladimir Volkov,† Thijs J. H. Vlugt,‡ and Theo W. de Loos*,‡ †
A.V. Topchiev Institute of Petrochemical Synthesis, Russian Academy of Sciences, Leninsky prospect 29, Moscow, 119991, Russian Federation ‡ Engineering Thermodynamics, Process & Energy Department, Faculty of Mechanical, Maritime and Materials Engineering, Delft University of Technology, Leeghwaterstraat 39, 2628 CB Delft, The Netherlands ABSTRACT: The solubility of CO2 in two biodegradable ionic liquids tris(2hydroxyethyl)methylammonium methylsulfate [THMA][MeSO4] and 2-hydroxyethyl-trimethylammonium lactate [2HETMA][Lac] has been studied experimentally. A synthetic method was used to measure bubble-point pressures up to 8 MPa for a temperature range of 313 K to 363 K. The solubility of CO2 in [2HETMA][Lac] is much higher than in [THMA][MeSO4], but the solubility increment is lower at higher CO2 concentrations. The solubility of CO2 in [THMA][MeSO4] is much lower compared to ionic liquids containing the same anion, but lacking hydroxyl-groups in the cation. The hydroxyl-groups required for the biodegradability of the IL have a detrimental effect on the CO2 solubility. [THMA][MeSO4] and [2HETMA][Lac] are not suitable for CO2 capture and can be considered as an example of a contradictive design where the biodegradability is improved upon introducing hydroxyl-groups, while the CO2 solubility is reduced significantly. The experimental data has been modeled with good accuracy using the Peng− Robinson equation of state in combination with the Wong−Sandler mixing rules.
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INTRODUCTION Carbon dioxide (CO2) is considered the main greenhouse gas, and is inevitably produced in large quantities by fossil-fuel based industries. Therefore, it is essential to limit the CO2 emissions into the atmosphere, and one of the tools to achieve this goal is carbon dioxide capture and storage (CCS).1 However, owing to the extremely large scale of the problem, the capture part of the CCS process turns out to be very challenging in terms of energy efficiency and cost.2 The currently used amine-based solvents for CO2 capture, at postcombustion conditions, have several drawbacks including their volatility, corrosivity, and high regeneration cost.2,3 To overcome some of these problems, ionic liquids (ILs) were proposed more than a decade ago as an alternative.4 ILs indeed bear properties (e.g., low vapor pressure, high stability, and high CO2 solubility) which make them suitable for CO2 capture, but their designation as “green solvents” is nowadays questioned from a green chemistry point of view.5 The main reason for this is that many ILs, especially the fluorinated ones, are toxic and nonbiodegradable.6 However, the tunability property of ILs can be used to make them biodegradable and nontoxic by including proper functional groups (e.g., hydroxyl and esters) in the anions/cations.7 Theoretically, improving the biodegradation should not affect the performance of the IL with respect to the CO2 solubility. This motivated us to select two relatively inexpensive and fully biodegradable8,9 hydroxylammonium-based ILs for CO2 solubility measurements and to compare the performance of © XXXX American Chemical Society
these ILs with the more commonly used ILs. Hence, we have measured the CO2 solubility in the IL tris(2-hydroxyethyl)methylammonium methylsulfate [THMA][MeSO4] and 2hydroxyethyl-trimethylammonium lactate [2HETMA][Lac] in a temperature range of 313 K to 363 K and up to 8 MPa. Henry’s constants have been extracted from which the enthalpy and entropy of absorption of CO2 are derived to assess the solute−solvent interactions. Subsequently, a comparison is provided of the experimental results and simulation results of the same CO2−IL system from the literature.
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EXPERIMENTAL SECTION The carbon dioxide used in the experiments had a purity of 99.995 mol % and was supplied by The Linde Group. The ILs tris(2-hydroxyethyl)methylammonium methylsulfate [THMA][MeSO4] and 2-hydroxyethyl-trimethylammonium lactate [2HETMA][Lac] (> 95 mol % purity) were purchased from Sigma-Aldrich. The ILs were dried prior to the experiments under moderate vacuum (p ≈ 0.01 mbar) at 80 °C for several days. The water content after drying was determined by Karl Fischer titration (Metrohm 756 KF Coulometer) and is shown together with other IL properties in Table 1. The first step in the sample preparation starts with filling a known amount of dried IL in a Pyrex-glass tube. This tube is then connected to a Received: August 13, 2013 Accepted: February 3, 2014
A
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Table 1. Properties of the ILs Used in this Worka
a
The water content was measured with Karl Fischer titration, whereas the melting point, purity, and biodegradability were reported by the supplier (Sigma-Aldrich).8.
⎡ ⎛a ⎞ A E (x ) ⎤ am = bm⎢∑ xi⎜ i ⎟ + ∞ i ⎥ ⎢⎣ i ⎝ bi ⎠ C ⎥⎦
gas dosing system, which is a part of our experimental procedure to degass the sample and to dose the gas. In the second step, the IL sample is degassed by freezing it with liquid nitrogen and applying high vacuum to the sample. The procedure of melting the sample, freezing it with liquid nitrogen, and evacuation is repeated until the sample is degassed completely (p < 0.0001 mbar) . In the third step, CO2 is dosed through a calibrated vessel into the tube using displacement by mercury, which also serves as a sealing and pressurizing fluid. The tube with a known composition is now disconnected from the gas dosing system and placed in the Cailletet apparatus. A brief description of the Cailletet apparatus will be provided here, since a detailed explanation can be found elsewhere.10 The Cailletet equipment operates according to the synthetic method and allows phase equilibria measurements within a pressure range of 0.1 MPa to 15 MPa and a temperature range of 255 K to 470 K. For a binary system, the phase rule requires two independent variables to be fixed at bubble-point conditions. Therefore, the composition of the CO2−IL mixture and the temperature were fixed in the experiments, while the pressure was gradually adjusted until the disappearance of the last gas-bubble, which corresponds to the bubble-point pressure. The solubility of CO2 then involves measuring bubble-point pressures at different, but fixed temperatures and compositions. The pressure is controlled by a dead weight gauge, while the temperature is measured by a Pt-100 thermometer. The uncertainty in the pressure, temperature, and composition of the measurements is ± 0.005 MPa, ± 0.01 K and ± 0.003 in the mole fraction, respectively.
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⎛ ⎛ a ⎞⎞ ⎛ ai ⎞⎞ ⎛ ⎛ aj ⎞⎞⎤ 1 ⎡⎛ ⎟⎟ = ⎟ ⎟ + ⎜b − ⎜ ⎟ ⎟⎥ ⎜b − ⎜ ⎢⎜bi − ⎜ j ⎝ RT ⎠⎠ ⎝ RT ⎠⎠ ⎝ ⎝ RT ⎠⎠⎦ ⎝ 2 ⎣⎝ ij (1 − kij)
GE = RT
τji =
(1)
∑i ∑j xixj(b − (a /RT ))ij E 1 − (∑i xiai /biRT ) − (A∞ (xi)/CRT )
∑ xi
∑ j = 1 τjiGjixj
i=1
gji − gii RT
N
∑l = 1 Glixl
(5)
, Gji = exp( −αjiτji), αji = αij
(6)
Table 2. Critical Temperature (Tc), Pressure (Pc), and Acentric Factor (ω) of the Components Used in the Modeling
where v is the molar volume, am and bm are constants for the mixture accounting for the molecular interaction and covolume, respectively. The PR EoS is used in combination with the Wong−Sandler (WS) mixing rules:12 bm =
N
N
where gji is the interaction energy between molecules i and j. The αji = αij parameter is related to the level of nonrandomness in the mixture. The PR EoS requires the critical properties and acentric factors of the components in the mixture in order to calculate the pure component parameters ai and bi.11 These properties are unknown for ILs, since many ILs decompose well before the critical point is reached. Here, the modified Lydersen−Joback−Reid group-contribution approach proposed by Valderrama et al.14 has been applied to calculate the critical properties of the investigated ILs. The critical properties and the acentric factor of the components are provided in Table 2.
The Peng−Robinson (PR) equation of state11 (EoS) has been used to model the experimental data, am RT − v − bm v(v + bm) + bm(v − bm)
(4)
where kij is a binary interaction parameter, ai and bi are pure component parameters accounting for the molecular attraction and covolume, respectively. C is an EoS dependent constant, which is −0.62322 for the Peng−Robinson EoS. AE∞(xi) is the excess Helmholtz energy at infinite pressure and is calculated from an appropriate activity coefficient model assuming AE∞(xi) ≈ AE0 (xi) ≈ GE0 (xi), which is the excess Gibbs energy at low pressure. The activity coefficient model used here is the nonrandom two-liquid (NRTL) equation,13
THERMODYNAMIC MODELING
P=
(3)
(2)
a
B
component
Tc (K)
Pc (MPa)
ω
[THMA][MeSO4]a [2HETMA][Lac]a CO2b
1256.6 806.0 304.12
0.820 2.594 7.374
0.831 1.325 0.225
Calculated from Valderrama et al.14 bTaken from Poling et al.30 dx.doi.org/10.1021/je400732q | J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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The bubble points of the CO2−IL systems have been calculated by simultaneously optimizing the parameters kij, αij = αji, τ12, and τ21 with a simplex algorithm in MATLAB,15 which minimizes the objective function: 1 OF = N
N
∑ i=1
Piexp − Pipred Piexp
Table 4. Bubble Point Data of the System CO2 (1) + [2HETMA][Lac] (2)a
(7)
Pexp i
Pprep i
where N the number of experimental data points, and are the experimental and predicted bubble-point pressures, respectively.
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RESULTS AND DISCUSSION The results of the CO2 solubility in [THMA][MeSO4] and [2HETMA][Lac] are provided in Table 3 and Table 4 and Table 3. Bubble Point Data of the System CO2 (1) + [THMA][MeSO4] (2)a x1
T/K
p/MPa
x1
T/K
p/MPa
0.030 0.030 0.030 0.030 0.030 0.030 0.059 0.059 0.059 0.059 0.059 0.059 0.114 0.114 0.114 0.114 0.114 0.114
313.23 323.18 333.17 343.24 353.26 363.28 313.22 323.20 333.18 343.20 353.22 363.25 313.23 323.18 333.17 343.34 353.27 363.20
0.679 0.814 0.959 1.115 1.280 1.460 1.388 1.648 1.953 2.278 2.630 3.015 3.848 4.453 5.163 5.952 6.775 7.629
0.045 0.045 0.045 0.045 0.045 0.045 0.087 0.087 0.087 0.087 0.087 0.087
313.25 323.24 333.21 343.24 353.27 363.26 313.20 323.24 333.21 343.24 353.27 363.25
1.034 1.240 1.465 1.690 1.965 2.215 2.149 2.605 3.100 3.641 4.241 4.802
x1
T/K
p/MPa
x1
T/K
p/MPa
0.042 0.042 0.042 0.042 0.042 0.082 0.082 0.082 0.082 0.082 0.171 0.171 0.171 0.171 0.171
323.47 333.37 343.46 353.49 363.47 323.47 333.53 343.55 353.56 363.60 323.28 333.25 343.22 353.27 363.29
0.200 0.290 0.401 0.531 0.686 0.592 0.837 1.102 1.422 1.778 2.335 3.086 3.851 4.742 5.748
0.063 0.063 0.063 0.063 0.063 0.120 0.120 0.120 0.120 0.120
323.22 333.22 343.25 353.26 363.29 323.19 333.22 343.22 353.23 363.25
0.370 0.531 0.721 0.941 1.196 1.164 1.570 2.040 2.580 3.161
a
x1 is the mole fraction of CO2, p is the bubble-point pressure, T is the temperature, and u(i) is the standard uncertainties: u(x1) = 0.003, u(T) = 0.01 K, and u(p) = 0.005 MPa.
a
x1 is the mole fraction of CO2, p is the bubble-point pressure, T is the temperature, and u(i) is the standard uncertainties: u(x1) = 0.003, u(T) = 0.01 K, and u(p) = 0.005 MPa.
Figure 1. Isopleths of the system CO2 + [THMA][MeSO4] for various carbon dioxide concentrations (mole %): 3.0 % (filled diamonds), 4.5 % (open squares), 5.9 % (filled triangles), 8.7 % (open triangles) and 11.4 % (filled circles). The lines through the data are polynomial fits to guide the eye.
graphically shown in Figure 1 and Figure 2, respectively. As expected the bubble-point pressure increases with increasing temperature for a fixed CO2 concentration. The solubility of CO2 in [THMA][MeSO4] is very low as can be observed in the P−x diagram shown in Figure 3. The bubble-point pressures increase sharply even at low CO2 mole fractions. The CO2 solubility in [2HETMA][Lac] is relatively high at low concentrations, but the bubble-point pressures for this system also increase sharply at higher CO2 concentrations as shown in Figure 4. Furthermore, Figure 3 and Figure 4 show that the PRWS/NRTL model, using the parameters listed in Table 5, can accurately represent the experimental data. Figure 5 provides a comparison of the CO2 solubility in the investigated ILs and three different ILs from the literature containing the same anion.16−18 Clearly, the CO2 solubility in [THMA][MeSO4] and [2HETMA][Lac] is significantly lower than in fluorinated ILs (e.g., [bmim][Tf2N]). More interesting is the comparison of the CO2 solubility in ILs from our previous work16 containing the same [MeSO4] anion and very similar cations, but without the hydroxyl groups in the cation. As can be seen in Figure 5, the CO2 solubility in the hydroxyl-containing IL (i.e.,
[THMA][MeSO4]) is much lower than in tributylmethylammonium methylsulfate [tbma][MeSO4], while the cation of both ILs are very similar, except that the [tbma] cation is lacking hydroxyl-groups. This suggests that the hydroxyl-groups have a detrimental effect on the CO2 solubility and contradict the generally accepted finding that cations play a minor role in CO2 absorption.19 This dramatic effect of hydroxyl-groups on CO2 solubility has also been observed by several other researchers.20−22 Aparicio et al.23,24 investigated the systems CO2 + [THMA][MeSO4] and CO2 + [2HETMA][Lac] in great detail by means of Molecular Dynamic (MD) simulations and the COSMO-RS approach. These authors explained the low solubility in [THMA][MeSO4] and [2HETMA][Lac] by the increased hydrogen bonding between the anions and hydroxyl-cations leading to a less effective rearrangements of the available cavities upon CO2 dissolution and consequently to low CO2 solubilities.24 In other words, the effective free volume C
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Figure 4. Isotherms of the system CO2 + [2HETMA][Lac] for several temperatures: 323.15 K (open squares), 333.15 K (filled triangles), 343.15 K (open circles), 353.15 K (open diamonds) and 363.15 K (filled circles). The lines are PR-WS/NRTL modeling results.
Figure 2. Isopleths of the system CO2 + [2HETMA][Lac] for various carbon dioxide concentrations (mole %): 4.2 % (filled diamonds), 6.3 % (open squares), 8.2 % (filled triangles), 12.0 % (open triangles) and 17.1 % (filled circles). The lines through the data are polynomial fits to guide the eye.
Table 5. EoS Parameters Obtained by Fitting the Calculated Bubble-Points to the Experimental Bubble-Points for the Systems CO2 (1) + [THMA][MeSO4] (2) and CO2 (1) + [2HETMA][Lac](2) T/K
Figure 3. Isotherms of the system CO2 + [THMA][MeSO4] for several temperatures: 313.15 K (filled diamonds), 323.15 K (open squares), 333.15 K (filled triangles), 343.15 K (open circles), 353.15 K (open diamonds) and 363.15 K (filled circles). The lines are PR-WS/ NRTL modeling results.
k12
313.15 323.15 333.15 343.15 353.15 363.15
0.277 0.451 0.406 0.599 0.721 0.721
323.15 333.15 343.15 353.15 363.15
0.046 0.029 0.031 −0.037 −0.039
α
τ12
CO2 + [THMA][MeSO4] 0.127 11.105 0.112 9.094 0.071 7.291 0.121 6.818 0.127 7.588 0.155 6.117 CO2 + [2HETMA][Lac] 0.045 43.763 0.052 30.319 0.060 23.833 0.073 18.770 0.107 22.078
τ21
AARD %
−0.965 −1.697 −2.671 −1.521 −1.543 −1.012
3.32 2.99 2.04 2.08 1.99 1.56
−6.242 −6.165 −5.399 −4.075 −1.073
4.53 4.68 4.76 4.92 6.05
is reduced due to the cation and anion interactions requiring increased pressures to host CO2 in the cavities of [THMA][MeSO4]. This molecular view of CO2 dissolution in the hydroxyl-containing IL is consistent with our experimental data as will be explained in the following. Determination of the Henry constant H12 of a solute 1 in a solvent 2 is often the first step in reducing gas-solubility data and can be obtained from
H12 = lim
x1→ 0
f1L x1
(8)
fL1
where x1 the CO2 mole fraction and the fugacity of the CO2 in the liquid. Since ILs have a very low vapor pressure25 it is justified to assume that the vapor phase is pure CO2 and to calculate fL1 from an equation of state26 of pure CO2 by applying the equilibrium condition:
f1L = f1V
Figure 5. Comparison of CO2 solubilities in [MeSO4] containing ILs at 353.15 K: [THMA][MeSO4] (crosses), [2HETMA][Lac] (plusses), [bmim][MeSO4] (squares), [tbma][MeSO4] (triangles), [tbmp][MeSO4] (circles), and [bmim][Tf2N] (diamonds).16 The lines are polynomial fits to guide the eye.
(9) D
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The Henry constants of CO2 in the ILs have been determined by plotting the fugacity of CO2 as a function of the CO2 mole fraction and taking the limit of x1 → 0. The results for different temperatures are provided in Table 6. The
ln[H12/MPa] =
CO2 + [2HETMA][Lac]
T/K
H12/MPa
T/K
H12/MPa
313.15 323.15 333.15 343.15 353.15 363.15
21.90 26.10 30.68 35.62 40.95 46.65
323.15 333.15 343.15 353.15 363.15
4.64 6.70 9.23 12.25 15.75
(10)
i=0
The coefficients ai of the BK equation are provided in Table 7. With the BK equation, the Gibbs energy, the enthalpy, and entropy of absorption of CO2 in the IL can be determined using the thermodynamic relations:29
Table 6. Henry Constants for CO2 (1) in the measured ILs (2) Obtained by Plotting the Fugacity versus the Mole Fraction and Taking the Limit of x1 → 0 CO2 + [THMA][MeSO4]
∑ ai(T /K)−i
ΔabsG∞ = RT[ln(H12)]
(11)
⎡ ∂ ⎛ Δ G ∞ ⎞⎤ ⎡ ∂ ln(H12) ⎤ ΔabsH ∞ = −T 2⎢ ⎜ abs ⎟⎥ = −RT 2⎢ ⎥ ⎣ ∂T ⎦ ⎣ ∂T ⎝ T ⎠⎦ (12) ∞
∞
(ΔabsH − ΔabsG ) T ⎡ ∂ ln(H12) ⎤ = −RT ⎢ ⎥ − R ln(H12) ⎣ ∂T ⎦
ΔabsS∞ =
Henry constants of CO2 in [THMA][MeSO4] are extremely high compared to the fluorinated [bmim][Tf2N] IL, which has a Henry constant of around 4.8 MPa at 60 °C.27 A comparison of the experimental and predicted Henry constants by the COSMO-RS approach from Aparicio et al.24 is provided in Figure 6. The predicted Henry constants for both the systems
(13)
The enthalpy and entropy of absorption can provide useful information on the dissolution behavior, since these properties are related to the strength of the interaction and the degree of ordering of the solvent molecules around a solute molecule, respectively. The calculated properties of the CO2 + [THMA][MeSO4], CO2 + [2HETMA][Lac], and the CO2 + [tbma][MeSO4] system from our previous work16 are provided in Table 7. The calculated enthalpy of absorption for the [2HETMA][Lac] system is relatively high suggesting the occurrence of a chemical reaction. This could explain the relatively good solubility of CO2 in [2HETMA][Lac] at low CO2 concentrations where the chemical absorption is the predominant phenomena. At higher CO2 concentrations, once the IL is saturated, the physical absorption phenomena becomes predominant, which also explains the sharp increase in the bubble-point pressures at the higher concentration scale. To infer the effect of the hydroxyl-groups on the dissolution properties, a comparison between two similar ILs, but one containing hydroxyl-groups (i.e., [THMA][MeSO4]) and the other lacking hydroxyl-groups (i.e., [tbma][MeSO4]) is given. CO2 is much more soluble in [tbma][MeSO4] than in [THMA][MeSO4], while the enthalpy of absorption of the hydroxyl-containing IL [THMA][MeSO4] is more negative (i.e., stronger interaction). The much lower CO2 solubility in [THMA][MeSO4] compared to [tbma][MeSO4] is due to entropic or free volume effects, which is consistent with the more negative value of the entropy of absorption for [THMA][MeSO4] and with the simulation results of Aparicio et al.24 The lower molar volume of [THMA][MeSO4] (i.e., 208.8 cm3/mol at 343.15 K) with respect to [tbma][MeSO4] (i.e., 302.8 cm3/mol at 343.15 K) is also consistent with the preceding conclusion. A comparison of Henry constants of CO2 in hydroxylammonium-based ILs with a lactate anion also
Figure 6. Comparison of the experimental Henry constants for (circles) [THMA][MeSO4] and (diamonds) [2HETMA][Lac] with the predicted Henry constants for (solid line) [THMA][MeSO4] and (dashed line) [2HETMA][Lac] by COSMO-RS.24
deviate significantly from the experimentally obtained Henry constants. The deviation for the [2HETMA][Lac] system is larger and becomes more pronounced at higher temperatures. The temperature dependency of the Henry constant is fitted to the Benson−Krause (BK) equation:28
Table 7. Coefficients of the BK Equation (eq 10), the Gibbs Energy (ΔabsG), the Enthalpy (ΔabsH) and Entropy of Absorption (ΔabsS)a ΔabsG
a
CO2 + IL
a0
a1
[THMA][MeSO4] [2HETMA][Lac] [tbma][MeSO4]b
10.938 14.943 6.528
−1738.231 −3582.003 −1295.08
ΔabsH
−1
kJ mol
16.76 12.91 14.43
−1
kJ mol
−14.45 −29.78 −10.77
ΔabsS J mol−1 K−1 −90.96 −124.41 −73.42
The values of ΔabsG, ΔabsH, and ΔabsS are consistent with a reference state of 0.1 MPa and 343.15 K. bCalculated from Ramdin et al.16 E
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shows the dramatic effect of the hydroxyl groups. For example, the Henry constants for CO2 in [2HETMA][Lac], 2-(2hydroxyethoxy)ammonium lactate [HEA][Lac],22 and bis(2hydroxyethyl)ammonium lactate [BHEA][Lac]20 are 3.07 MPa, 18.67 MPa, and 22.2 MPa at 313.15 K, respectively. Clearly, the CO2 solubility decreases if the number of hydroxyl-groups in the cation increases from 1 for [2HETMA][Lac] to 2 for [HEA][Lac] and [BHEA][Lac]. The main conclusion is that the CO2 interacts more strongly with the hydroxyl-containing IL, but at the expense of the free volume. Both investigated ionic liquids are not suitable for CO2 capture and can be considered as an example of a contradictive design where the biodegradability is improved upon introducing hydroxylgroups, while the CO2 solubility is reduced significantly. Hence, groups which promote hydrogen bonding between the ions of the IL reduce the free volume and should be eliminated at the functionalization step if one aims to improve the CO2 solubility and biodegradability at the same time.
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CONCLUSIONS The solubility of CO2 in the biodegradable ILs tris(2hydroxyethyl)methylammonium methylsulfate [THMA] [MeSO4] and 2-hydroxyethyl-trimethylammonium lactate [2HETMA][Lac] has been measured using a synthetic method. The CO2 solubility in [THMA][MeSO4] and [2HETMA][Lac] is low compared to other commonly used ILs. This low solubility is caused by the strong cation and anion interactions thereby reducing the free volume and hence also the solubility. A comparison of the CO2 solubility in [THMA][MeSO4] with an IL having the same anion and a similar cation, but lacking the hydroxy-groups in the cation, shows that the hydroxylgroups have a detrimental effect on the solubility. [THMA][MeSO4] and [2HETMA][Lac] can be considered as an example of a contradictive design where the biodegradability is improved upon introducing hydroxyl-groups, while the CO2 solubility is reduced significantly. Groups which promote hydrogen bonding between the ions of the IL reduce the free volume and should be eliminated at the functionalization step if one aims to improve the CO2 solubility and biodegradability at the same time.
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Funding
Financial support by the Advanced Dutch Energy Materials program of the Dutch Ministry of Economic Affairs, Agriculture and Innovation is acknowledged. This research cooperation is also a part of Dutch-Russian Centre of Excellence “Gas4S” (NWO-RFBR No. 047.018.2006.014/08-08-92890-CE). Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS The authors thank Ir. E. J. M. Straver for his assistance with the experimental work.
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REFERENCES
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