Absorption Behavior of Acid Gases in Protic Ionic ... - ACS Publications

Mar 7, 2017 - Nizhny Novgorod State Technical University n.a. R.E. Alekseev, 24 Minina ... N.I. Lobachevsky State University of Nizhny Novgorod, 23 Ga...
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Research Article pubs.acs.org/journal/ascecg

Absorption Behavior of Acid Gases in Protic Ionic Liquid/ Alkanolamine Binary Mixtures Alsu I. Akhmetshina,† Anton N. Petukhov,† Andrey V. Vorotyntsev,† Alexander V. Nyuchev,‡ and Ilya V. Vorotyntsev*,† †

Nizhny Novgorod State Technical University n.a. R.E. Alekseev, 24 Minina str., Nizhny Novgorod 603950, Russian Federation N.I. Lobachevsky State University of Nizhny Novgorod, 23 Gagarin Avenue, Nizhny Novgorod 603950, Russian Federation

‡

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S Supporting Information *

ABSTRACT: Herein, we studied the absorption of H2S and CO2 by alkanolamine−protic ionic liquids binary mixtures based on 2-hydroxyethylammonium (MEA) or triethanolammonium cations and residues of 2hydroxy-5-sulfobenzoic acid or pyridine-3-carboxylic acid at various temperatures and partial gases pressures. It was found that absorbents based on the 2-hydroxyethylammonium cation, performed high absorption properties toward the H2S. The solubility of hydrogen sulfide, characterized by the Henry’s Law constant, in MEA-based binary mixtures had the values comparable to the commercially available ionic liquids. The results of thermal desorption analysis demonstrated that the capture of acid gases in MEAbased absorbents occurred at two stages: through the dissolution in MEA component and in protic ionic liquid. KEYWORDS: Protic ionic liquids, Alkanolamines, Acidic gases, Solubility, Thermodynamics

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INTRODUCTION The typical absorbents for acidic gases are ethanolamines combining the alkalinity of amine group, which allows them interacting with acidic compounds and the ability of hydroxyl groups to reduce the vapor pressure. Industrially applied forms of amine-based absorbents are centered on aqueous solutions of ethanolamines. In general, the chemical absorption of acidic gases in aforementioned solutions occurs by the direct reaction of CO2 with amine to form the carbamate ions. Selection of primary, secondary or tertiary ethanolamine depends on the treated gas composition.1,2 Moreover, the affinity toward the acidic gases is determined by the basicity and pKa values of amines increasing from primary to tertiary3,4 Also, the selectivity of absorption varies by the ethanolamines chemical structure. It is well established that primary ethanolamines are highly effective for the capture of all acidic gases (hydrogen sulfide, carbon dioxide and carbonyl sulfide). Secondary ethanolamines, predominantly, react with CO2 and H2S and much less with COS, whereas tertiary ethanolamines provide higher H2S removal capacity than other acidic gases. However, some drawbacks of amine-based technique such as amine solutions degradation, emission of volatile organic compounds5,6 and corrosion activity became premises toward the research of novel routes for acidic gases removal. Numerous works mentioned potential alkanolamines replacement by the various types of ionic liquids (ILs), which are salts, composed of bulky organic cations and organic or inorganic anions melting below 373 K. Extremely low vapor pressure, high thermal stability over a wide range of © 2017 American Chemical Society

temperatures and high values of sorption capacity make ILs the promising candidates to remove acidic gases from the gas steams.7 Several studies reported the solubility of CO2 in imidazolium-based ILs, which absorbed acid gases physically at high pressures.8−17 However, at the ambient pressure conventional ILs perform relatively low values of gases solubility. The enhanced sorption capacity can be achieved in task-specific ILs containing the additional amine groups in cation or anion. Despite the high sorption capacity, amine-functionalized ionic liquids had a high viscosity limited their application in industrial scales.18 Interesting subset of ILs known as protic ionic liquids is produced through the proton transfer from the Brønsted acid to the Brønsted base.19 Several protic ILs with alkanolaminebased cations were investigated in order to absorb acidic contaminates from gas mixtures. A series of protic ILs with hydroxyl group containing ammonium cation were trialed in work of Yuan et al.20 for CO2 capture, and the highest solubility values were found for tri-(2-hydroxy ethyl)-ammonium lactate and 2-(2-hydroxy ethoxy)-ammonium acetate. Protic ILs obtained from N-methyl-2-hydroxyethylammonium paired with formate or acetate anions have performed high CO2 solubility under the elevated pressures of up to 80 MPa.21 As an alternative, several works based on the combination of ethanolamines and ionic liquids for the prevention of amine Received: January 10, 2017 Revised: February 10, 2017 Published: March 7, 2017 3429

DOI: 10.1021/acssuschemeng.7b00092 ACS Sustainable Chem. Eng. 2017, 5, 3429−3437

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been measured in temperature range from 293 to 333 K, and activation energy of viscous flow was calculated.

losses and diminishing the energy consumption have been reported as novel absorbents. This type of absorbents provides the main advantages, such as absorption capability of mixtures for acidic gases comparable to that of amines solutions, relatively low cost and low partial pressures for acceptable gases capture.22,23 It was the Prof. Noble’s (University of Colorado Boulder, USA) and his research group’s idea to combine the ILs and amines for rapid and reversible gases removal.24 It was found that the mixtures of ILs and amines absorb 1 mol of CO2 per 2 mol of dissolved amine.25 The approach of mixing of ILs with amines also was highlighted in work of Yang,22 who reported that the mixed IL − amine solution (30 wt % MEA + 40 wt % [bmim][BF4] + 30 wt % H2O) exhibited slightly higher CO2 removal efficiency than the corresponding 30 wt % MEA aqueous solution. Yu with co-workers23 investigated the CO2 absorption capability of binary mixtures of 1-butyl-3methylimidazolium hexafluorophosphate−monoethanolamine, 1-ethyl-3-methylimidazolium tetrafluoroborate−monoethanolamine, 1-butyl-3-methylimidazolium tetrafluoroborate−monoethanolamine and 1-ethyl-3-methylimidazolium hexafluorophosphate−monoethanolamine, and established that the CO2 absorption capacity in ILs−MEA compositions became much higher than that in the neat ILs and was comparable to the values in some aqueous amines and mixture of ILs−amine− H2O. Ahmady et al. investigated aqueous solution of methyldiethanolamine mixed with three types of ionic liquids: 1-butyl-3-methyl-imidazolium tetrafluoroborate, 1-butyl-3methyl-imidazolium acetate and 1-butyl-3-methyl-imidazolium dicyanamide.26 Osman prepared hybrid solvents on the basis of [bmim][BF4] combined with MEA, DEA, and MDEA with different concentrations, that achieved superior CO2 absorption over alkanolamines.27 The protic ionic liquids were suggested in a range of reports as the acid gases absorbents with a low cost. An important feature thereof is the fewer amounts of basic amine groups in protic ILs than in amines, leading to the lower acidic gases capture ability and lower energies required for the desorption of gases.19 The primary amines and, in particular, monoethanolamine, represent the strongest alkaline properties. Therefore, we considered the monoethanolamine as a potential candidate to design protic ILs with the high absorption capacity. The tertiary amines are less alkaline, however, are able to selective absorption of hydrogen sulfide due to the different rates of reaction with H2S and CO2. Among the tertiary amines, the triethanolamine was chosen as the most commercially available and widely used to design the H2S-selective protic ionic liquids. In the literature, ionic liquids with remarkable biodegradability containing toxicologically inert salicylate28 and nicotinic acid derivatives29 were proposed for the CO2 adsorption,30 as a reaction media and pH-responsible materials.31 In this work, we selected anions of nicotinic acid and 5-sulfosalicylic acid, namely, pyridine-3-carboxylate and 2-hydroxy-5-sulfobenzoate, to prepare the protic ionic liquids for acidic gases removal. Seeing this impressive precedent, this work has inspired the current study toward the experimental determination of CO2 and H2S solubility in amine mixtures with the four carboxylatebased protic ionic liquids 2-hydroxyethylammonium pyridine-3carboxylate, 2-hydroxyethylammonium 2-hydroxy-5-sulfobenzoate, triethanolammonium (TEA) pyridine-3-carboxylate, triethanolammonium 2-hydroxy-5-sulfobenzoate at different temperatures and partial acidic gases (CO2 and H2S) pressures. Moreover, the density and viscosity of the binary mixtures have

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EXPERIMENTAL SECTION

Materials. the reagents ethanolamine (≥98% purity), triethanolamine (≥98% purity), 2-hydroxy-5-sulfobenzoic acid (Sulf) dehydrate (≥99% purity) and pyridine-3-carboxylic acid (Nic) (98% purity) were purchased from Sigma-Aldrich Ltd. and used without any further purification. The carbon dioxide and hydrogen sulfide used were delivered from “Monitoring” (Russia) with a purity of 99.95% and the helium purity was more 99.99999%. General Procedure for Preparation of Absorbents. The protic ionic liquids 2-hydroxyethylammonium pyridine-3-carboxylate (MEA[Nic]), 2-hydroxyethylammonium 2-hydroxy-5-sulfobenzoate (MEA[Sulf]), triethanolammonium pyridine-3-carboxylate (TEA[Nic]) and triethanolammonium 2-hydroxy-5-sulfobenzoate (TEA[Sulf]) were prepared via neutralization reaction between the amine and the acid. In brief, the calculated amount of acid solution in ethanol was introduced in a three-necked flask, which was equipped with a magnetic stirrer, thermometer and a drop funnel. Then, the equimolar amount of base was added dropwisely over a period of 1 h. The reaction mixture was stirred for 5 h at 273.15 K. Next, the excess amount of alkanolamine equal to 1 mol was added to the reaction mixture. After the procedure, the solvent was removed via rotary evaporation by RV 10 digital V (IKA, Germany) under reduced pressure for 12 h. All absorbents were weighted by an analytical balance AUW-220D (Shimadzu, Japan), and the accuracy was ±0.0001 g. The water content in absorbents was determined by the Karl Fischer method by 831 KF Coulometer (Metrohm AG, Switzerland), and was equal to 3.49, 6.21, 3.39, 4.13% w/w MEA[Nic]−MEA, 2-aminoethanol 2-hydroxy-5-sulfobenzoate MEA[Sulf]−MEA, TEA[Nic]− TEA, TEA[Sulf]−TEA binary mixtures, respectively. IR and NMR Spectroscopy. The synthesized protic ILs structure and purity were identified by IR, 1H and 13C NMR spectroscopy. IR spectra were recorded at ambient temperature using FTIR spectrometry (IRAfinity-1, Shimadzu, Japan). A minimum of 30 scans was signal-averaged with resolution of 4 cm−1 at the 2000−800 cm −1 range. All other parameters were not controlled and corresponded to the testing characteristics established by the producer. The accuracy of measurement of wavenumbers was monitored by the spectrum of polystyrene, being ±0.2 cm−1. The sample measurements were carried out using the following method: samples were treated in a potassium bromide matrix according to the literature.32−35 1H NMR and 13C NMR spectra of protic ILs were recorded on an DD2 400 (Agilent, USA) spectrometer. Chemical shifts (δ) were reported in ppm for the solution of compound in DMSO-d6 with internal reference TMS and J values in hertz. Density and Viscosity Measurements. Densities and viscosities were measured using a SVM 3000 Stabinger Viscometer (Anton Paar, Austria) with an error of 0.000 05 g cm−3. The instrument was calibrated using water at 298.15 K. After calibration, the densities of MEA and TEA were measured at 298.15. The obtained values were compared with the literature data, with an average absolute deviation of 0.06%. The measurements were investigated at temperatures of 293, 303, 313, 323 and 333 K at constant atmospheric pressure. Each measurement point was repeated three times to ensure reproducibility. Thermal Desorption Analysis. For coupled evolved gas-analysis (Direct EGA-MS) measurements, a temperature-programmable MultiShot Pyrolyzer EGA/PY-3030D (Frontier Laboratories, Japan) incorporated in CGMS QP-2010Plus (Shimadzu, Japan) was used. To perform an experiment, a deactivated stainless steel sample cup loaded with approximately 50 μg of sample was dropped into a quartz pyrolysis tube. The quartz tube was surrounded by a tubular furnace, which provided uniform heating and maintained the pyrolysis temperature in accordance with the following temperature program under the helium flow at 50 mL/min, at the pyrolyzer was reduced to 1 mL/min at the capillary column by means of splitter: first stage, holding at 323.15 K for 10 min; after that, at the second stage the sample was heating in the range of 323.15−773.15 K (10 K/min). 3430

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During the second stage, identification of resulting chromatographic peaks was carried out with the help of a mass selective detector at GCMS QP-2010Plus (Shimadzu, Japan). Ionization in the mass spectrometer was carried out by electron impact (EI) at 70 eV and a mass range between 12 and 500 amu scanned at a rate of 2000 s/scan. Reaction products were separated on 2.5 m tube made of Ultra ALLOY EGA (Frontier Laboratories, Japan) at 373.15 K (45 min) with the help of selected ion monitoring, carrier gas was helium with purity 99.99999%. Reaction products were identified with the help of NIST-11 database of mass spectra and GC−MS real time analysis software. Determination of Gases Solubility. The experimental set up is schematically represented in Figure 1, and based on the pressure drop

Research Article

RESULTS AND DISCUSSION

Characterization of Protic Ionic Liquids and Their Mixtures with Alkanolamines. Absorbents on the basis of protic ionic liquid and alkanolamines were prepared in one step via the neutralization reaction of carboxylic acid and alkanolamine with 1 mol excess of the latter as shown in Figure 2.

Figure 2. Formation of 2-hydroxyethylammonium 2-hydroxy-5sulfobenzoate (a), triethanolammonium 2-hydroxy-5-sulfobenzoate (b), 2-hydroxyethylammonium pyridine-3-carboxylate (c), triethanolammonium pyridine-3-carboxylate (d). Figure 1. Scheme of the experimental setup: 1, samples; 2, vacuum pump post; 3, gas container; 4, high pressure equilibrium cell; 6, pressure sensor; 7, thermostatic air bath; 8, 9, temperature sensor; 10, magnetic stirring system; 11−13, valves; 14, gas cylinder.

Involvement of reagents into the neutralization reaction and the estimation of the unreacted reagents were conducted using IR spectra analysis. To establish the ammonium salt formation, we observed the bands corresponding to carboxylate and ammonium structures. The ammonium structures presence was determined by stretching vibrations of NH bonding centered at nearby 1600 cm−1 and stretching vibrations of CO bond in carboxylates within the 1300−1400 and 1550−1650 cm−1 range. Formation of the latter structures was accompanied simultaneously by disappearing of carbonyl stretching vibrations in carboxylic group at 1680−1700 cm−1.32 The complete formation of pyridine-3-carboxylates in binary mixtures was determined according to the presence of symmetric and asymmetric stretching vibrations of CO bond in carboxylate anions in the range of 1382−1385 cm−1 and 1601−1605 cm−1 (Figure 3), respectively, and the absence of stretching vibrations of CO bond in carboxylic groups in the IR spectrum.32 The peak corresponding to the plane vibrations of NH bonding in 2-aminoethanol at 1627 cm−1 shifts to lower frequency (1608 cm−1) in 2-hydroxyethylammonium pyridine-3-carboxylate. The similar tendency was observed in 2-hydroxy-5-sulfobenzoate-based salts, where the carbonyl stretching vibrations of 2-hydroxy-5-sulfobenzoic acid at 1670 cm−1 transformed to the stretching vibrations of CO bond in carboxylate in the range of 1630−1632 cm−1 and 1332−1337 cm−1 (Figure 4). The strong band at 1031 cm−1 implies formation of sulfonates. The additional information about the structure of alkanolamine−protic IL mixtures was obtained using 1H and 13C NMR spectroscopy (Supporting Information). Because of the fast exchange of the proton, it is impossible to distinguish the amine form from the protonated

method that is used in most experimental studies in the literature and described in publications.36−39 The apparatus consisted of a high pressure equilibrium cell (HPEC) (4) equipped with a magnetic bar of 16.72 cm3 volume, and a gas container (GC) (3) for the introduction of known amounts of 35.24 cm3 CO2. The volumes of different compartments of the setup were measured and precalibrated using a calibrated bulb of the known volume. The pressures in GC and HPEC were monitored using two pressure transducers CPG 1000 (5, 6) (WIKA GmbH, Germany) with an accuracy of ±0.05% FS positive pressure, ± 0.25% FS vacuum/500 psi and below). The temperature was maintained constant by a thermostatic air bath (7) applied with temperature sensor (8, 9) with an accuracy of ±0.1 K. The preweighted amount of PILs (1−2) was introduced into the equilibrium cell, and the system was evacuated with valves 11 and 13 open. The PILs were degassed and dried during 24 h under vacuum at 343.15 K. Then the valves were closed and the certain amount of gas (H2S or CO2) was introduced into the GC (3) from the cylinder with pure gas (H2S or CO2) through valve 11. Afterward, the valve to HPEC (12) was opened so that PILs could be in contact with gas. The determination of solubility at different temperatures was simply done by changing the air thermostat set point and waiting for a new thermodynamic equilibrium. With the single loading it was thus possible to make measurements over the large temperature range (303.15 to 333.15 K) in this study. Absorption equilibrium was thought to be reached when the pressures of the HPEC remained constant for at least 2 h. The difference between two PVT measurements was then used to calculate the quantity of solute present in PILs. 3431

DOI: 10.1021/acssuschemeng.7b00092 ACS Sustainable Chem. Eng. 2017, 5, 3429−3437

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Figure 3. FTIR spectra of TEA[Nic]−TEA (a), MEA[Nic]−MEA (b) mixtures.

Figure 5. Viscosity of MEA[Sulf]−MEA (1), TEA[Sulf]−TEA (2), TEA[Nic]−TEA (3), MEA[Nic]−MEA (4), bmim[PF6] (5), bmim[BF4] (6), bmim[Tf2N] (7).

Table 1. Activation Energy for ILs, Calculated by the Guzman-Andrade Equation

Figure 4. FTIR spectra of TEA[Sulf]−TEA (c), MEA[Sulf]−MEA (d) mixture.

η0 × 106, mPa·s

E, kJ·mol−1

correlation coefficient, R2

bmim [PF6] bmim [BF4] bmim [Tf2N] TEA[Sulf]:TEA = 1:1 TEA[Nic]:TEA = 1:1 MEA[Sulf]:MEA = 1:1 MEA[Nic]:MEA = 1:1

1.50 222.70 444.10 1.91 0.37 0.92 5.33

45.97 32.29 28.87 53.23 56.62 51.14 46.29

0.9980 0.9956 0.9982 0.9993 0.9987 0.9992 0.9979

comparison with the obtained absorbents, the E was higher than in the case of conventional ILs. The physicochemical properties of mixtures vary from the cation and anion structure. Alkanolamine−protic IL mixture with [Nic] anion displays lower activation energy values than salts based on the [Sulf]. This observation can be explained by the presence of two proton-donating groups in 2-hydroxy-5-sulfobenzoic acid instead of the one carboxylic group in pyridine-3-carboxylate. The hydroxyl groups of ethanolamines are also able to increase the energy barrier providing the additional hydrogen bonding, indeed, in current work we observed the correlation between the number of OH groups and the energy value. TEA containing absorbents with three hydroxyl groups in cation exhibit higher values of viscosity and activation energy than the MEA-based mixtures with one OH-group. The temperature dependences of protic ILs density have the linear behavior and are represented in the Figure 6. The densities of mixtures containing [Sulf] anion are higher than those with [Nic] anion. Also, it was noted that the cation had a subsequent impact to the density value. In particular, protic ILs with TEA cation exhibit higher density than MEA containing ILs. Acidic Gases Solubility. Throughout the literature, protic ionic liquids have shown high sorption properties19 toward the gases with active protons in molecule, such as hydrogen sulfide, interacting with electron-donating hydroxyl or carboxylate groups in protic ILs. The interaction of CO2 and protic ionic liquids is weaker than that of the aforementioned gases and occurs owing to the several mechanisms, such as Lewis acid− base interactions and the intermolecular interactions (electrostatic, hydrogen bonding, and van der Waals).46 All acidic gases

amine in the NMR spectra. Therefore, in 1H NMR spectra, the amine and protonated amine peaks exhibit the broad singlet in the 5.25−5.90 ppm range, which represents the OH group as well. By comparison, the viscosity data of bmim[PF6], bmim[BF4] and bmim[Tf2N] was observed in the literature,40−42 and the results are depicted in Figure 5. The experimental viscosity data was analyzed in order to determine the activation energy and fitted as a function of temperature, using the Guzman-Andrade equation:43

η = η0e E / RT

ionic liquid

(1)

where η is viscosity, mPa·s; E is activation energy of viscous flow, kJ·mol−1; T is temperature, K; R is the universal gas law constant, J·K−1·mol−1; η0 is viscosity at infinite temperature, mPa·s. The activation energy for viscous flow E and the viscosity at infinite temperature η0 were calculated from the slope of the plot (Table 1). The activation energy indicates the energy needed to overcome the intermolecular interactions between the molecules to change their positions. The higher values of activation energy correspond to the more difficult detachment of ion pairs, which is observed due to the strong intermolecular interactions or the presence of steric hindrances. According to the literature data, the activation energy values for conventional ILs bmim[PF6], bmim[BF4] and bmim[Tf2N] are 43.47,43 34.044 and 26.0545 kJ/mol, respectively. In 3432

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Table 3. Thermodynamic Properties of H2S in Binary Mixtures

Figure 6. Density of MEA[Sulf]−MEA (1), TEA[Sulf]−TEA (2), TEA[Nic]−TEA (3), MEA[Nic]−MEA (4).

dissolve in protic ILs through the physical sorption mechanism according to Henry’s law. Moreover, the solubility of acidic gases in amines is based on chemisorption of 2 mol of CO2 per 1 mol of amine. The results of CO2 and H2S solubility measurements at temperatures of 303.2, 313.2, 323.2, 333.2 K and pressures of up to 12.7 bar are summarized in the Supporting Information. The reliability and accuracy of the measurement method have been checked using [bmim][Tf2N],40,41 the average reproducibility of the solubility data was well within ±1%. The H(0) 21.m values at each temperature T can be obtained from the intercept of Krichevsky− Kasarnovsky plots.47 The molality-scale Henry’s law constants at zero pressure, H(0) 21.m are given in Tables 2 and 3 for the solubility of CO2 and H2S in the binary mixtures. The relative error associated with the measured Henry’s law constants is approximately 1.3%. The Henry’s law constants of H2S and CO2 in four absorbents at 303.2 K vary from 1.52 to 12.70 bar, with the

H(0) 21.m, bar

303.2 313.2 323.2 333.2

22.9087 26.3186 29.9342 36.7964

303.2 313.2 323.2 333.2

39.2377 46.0053 52.2423 100.7690

303.2 313.2 323.2 333.2

21.4691 27.2512 33.8058 50.3464

303.2 313.2 323.2 333.2

37.4584 41.0956 44.8760 51.6941

ΔsolG, kJ·mol−1

ΔsolH, kJ·mol−1

MEA[Sulf]−MEA 7.8901 −11.6546 8.5115 −13.0882 9.1290 −14.6318 9.9829 −16.2893 TEA[Sulf]−TEA 9.2459 −22.4691 9.9651 −25.2798 10.6247 −28.3087 12.7724 −31.5638 MEA[Nic]−MEA 7.7266 −21.3231 8.6021 −24.0571 9.4556 −27.0067 10.8510 −30.1799 TEA[Nic]−TEA 9.1291 −8.2719 9.6713 −9.1526 10.215 −10.0936 10.924 −11.0969

H(0) 21.m, bar

303.2 313.2 323.2 333.2

12.6236 14.4680 16.0217 17.5810

303.2 313.2 323.2 333.2

18.5234 21.4205 24.4943 25.0503

303.2 313.2 323.2 333.2

13.7323 15.2919 17.9561 22.5197

303.2 313.2 323.2 333.2

23.6146 28.2291 30.0072 34.5113

ΔsolG, kJ·mol−1

ΔsolH, kJ·mol−1

MEA[Sulf]−MEA 6.3885 −7.4805 6.9542 −8.4711 7.4502 −9.5415 7.9379 −10.6947 TEA[Sulf]−TEA 7.3547 −5.2161 7.9755 −5.8919 8.5903 −6.6214 8.9182 −7.4065 MEA[Nic]−MEA 6.6007 −13.0786 7.0983 −14.7707 7.7563 −16.5971 8.6234 −18.5626 TEA[Nic]−TEA 7.9666 −10.07805 8.6939 −11.2428 9.1355 −12.4931 9.8054 −13.8317

ΔsolS, J·mol−1·K−1 −45.7424 −49.2509 −52.5737 −55.9202 −41.4609 −44.2769 −47.0662 −48.9942 −64.9055 −69.8248 −75.3510 −81.5906 −59.5141 −63.6551 −66.9204 −61.5251

lowest for TEA[Nic]−TEA, and the largest value of this parameter was observed for MEA[Sulf]−MEA. However, the Henry’s law constants of CO2 are of a magnitude larger than that of H2S, implying that the binary mixtures exhibit larger absorption capacity for H2S than for CO2. Therefore, these protic ILs enable selective separation of H2S from CO2. It is noted that the solubility of CO2 in two Nic-based mixtures is much lower than that in Sulf-containing absorbents. One possible explanation for such result is the difference in structures of protic and aprotic ILs. In protic ILs, the protonated nitrogen in ammonium cation is weakly acidic. It can form Brønsted acid−base interaction (or hydrogen bond) with the alkaline carboxylates, thus reducing the CO2 affinity for those anions. However, in aprotic ILs, such interaction does not exist and the carboxylate anion can attract CO2 freely through Lewis acid−base interaction.48 Hydrogen sulfide is more favorable to be absorbed by protic ILs than CO2 owing to the existence of active protons in H2S molecule. As a result, H2S molecule still has the possibility of interacting with the electron-donating groups in protic ILs (e.g., hydroxyl and carboxylate groups). Theoretical calculations were carried out to demonstrate the different absorption behavior of H2S and CO2 in PILs-containing absorbents. Subsequently, the changes of Gibbs free energy and entropy of H2S and CO2 absorption in the PILs at different temperatures can be calculated from the following two equations. Thermodynamic properties of solution (of CO2 in absorbents) can be calculated from the correlation of Henry’s constant given above applying well-known thermodynamic relations:47

Table 2. Thermodynamic Properties of CO2 in Binary Mixture T, K

T, K

ΔsolS, J·mol−1·K−1 −64.4617 −68.9648 −73.5174 −78.8484 −104.6009 −112.5316 −120.4625 −133.0619 −95.8104 −104.2760 −112.8168 −123.1422 −57.3912 −60.1021 −62.8408 −66.0901 3433

Δsol G = RT ln(H21, m(T , p)/p°)

(2)

⎛ ∂ln(H21, m(T , p)/p°) ⎞ ⎟ Δsol H = R ⎜ ∂(1/T ) ⎠p ⎝

(3)

Δsol S = (Δsol H − Δsol G)/T

(4)

DOI: 10.1021/acssuschemeng.7b00092 ACS Sustainable Chem. Eng. 2017, 5, 3429−3437

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ACS Sustainable Chemistry & Engineering where ΔsolG, ΔsolH and ΔsolS are the Gibbs free energy of solvation, enthalpy and entropy of solvation, respectively. At standard ambient temperature and pressure (T° = 298.15 K, p° = 1 bar) the thermodynamic properties of CO2 and H2S in studied absorbents are calculated. The ΔsolH values and ΔsolS have negative values. The ΔsolH values and ΔsolS values are negative for all absorbents at any temperature, and they decrease when it rises. [TEA][Nic]−TEA+H2S and [TEA][Nic]−TEA+CO2 systems are found to have the similar values of ΔsolS (−57.3912 and −59.5141 J·mol−1·K−1), although the absorption of H2S differs largely in capacity from that of CO2 (H(0) 21.m= 37.4584 vs 23.6146 bar at 303.2 K). However, the absorption of H2S in the [TEA][Nic]−TEA has less negative values of ΔsolH than that of CO2. Therefore, it is concluded that the effect of enthalpy rather than entropy contributes primarily to the different absorption behavior of H2S and CO2 in the [TEA][Nic]−TEA. Because more negative ΔsolH leads to smaller ΔsolG in the absorption, the dissolution of H2S in two [MEA]-based protic ILs is thermodynamically more favorable than that of CO2. To select the appropriate absorbent for H2S capture and to compare the absorption capacity of the binary mixtures, conventional and protic ILs, we summarized the experimental results and literature data in Table 4. The Henry’s Law constant Table 4. Henry’s Law Constants for Alkanolamine−Protic IL Binary Mixtures and Conventional ILs at 303.2 K absorbent

H(0) 21.m, bar

[MDEAH][Ac] [MDEAH][For] [DMEAH][Ac] [DMEAH][For] emim[Tf2N] bmim[Tf2N] hmim[Tf2N] bmim[PF6] emim[eFAP] hemim[BF4] emim[EtSO4] omim[PF6] MEA[Sulf]−MEA TEA[Sulf]−TEA MEA[Nic]−MEA TEA[Nic]−TEA

5.5 11.5 3.5 5.9 14.8 13.7 17.4 27.7 15.3 31.3 60.7 64.4 12.6 18.5 13.7 23.6

ref. 45 45 45 45 46 47 48 49 50 51 52 53 This This This This

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at the ambient temperature was the parameter characterizing the absorption capacity. In addition, the mole fraction solubility of the studied absorbents was compared with the data known for protic ionic liquids and included in the Table 4. Thermal Properties. Protic ionic liquids are commonly produced by proton transfer between the acid and the base. The effect of proton transfer on the cation results in growth of desorption temperature and, consequently, the volatility of absorbent, as it was reported previously by Vijayraghavan and co-workers.50 The thermal desorption (TD) analysis was used to assess the volatility of protic ILs-etanolamine mixtures. Figure 7 shows the thermal degradation profiles of absorbents (a) TEA[Sulf]− TEA, (b) MEA[Sulf]−MEA, (c) TEA[Nic]−TEA, (d) MEA[Nic]−MEA. The samples were observed by Direct-EGA-MS using selected ion monitoring (SIM) modes. In Figure 7, the

Figure 7. EGA profiles of selected ion monitoring curves measured by Direct EGA-MS: (a) TEA[Sulf]−TEA, (b) MEA[Sulf]−MEA, (c) TEA[Nic]−TEA, (d) MEA[Nic]−MEA.

SIM curves at different m/z are represented by asymmetrical peaks. The degradation temperature was determined by the increase in intensity m/z 44 of carbon dioxide together with the increase 3434

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ACS Sustainable Chemistry & Engineering Table 5. Thermal Stability of Absorbents and Temperature Range of Degradation absorbent

m/z

compound name

temperature range, K

Tda, K

TEA[Sulf]−TEA

18 44 64 94 100 118 18 30 44 64 94 18 44 79 118 123 18 30 44 79 123

hydrogen dioxide carbon dioxide aminomethanesulfonic acid phenol 4-morpholineethanesulfonic acid TEA hydrogen dioxide MEA carbon dioxide aminomethanesulfonic acid phenol hydrogen dioxide carbon dioxide pyridine TEA nicotinic acid hydrogen dioxide MEA carbon dioxide pyridine nicotinic acid

323−673 423−623 523−723 523−723 473−623 373−573 423−573 373−523 323−573 473−623 423−623 323−573 423−573 473−573 373−573 423−523

>508

MEA[Sulf]−MEA

TEA[Nic]−TEA

MEA[Nic]−MEA

a

>453

>473

>433 323−503 323−493 373−523

Degradation temperature.

in intensity m/z 94, 64, 100 (for Sulf-containing protic ILs), m/ z 64, 94, 79 (for Nic-containing protic ILs). According to the results of TD analysis (Table 5), TEA[Sulf]−TEA absorbent performed the highest stability at the elevated temperatures. Therefore, the selection of TEA as the cation enhances the stability of an average at 50°. It is worth to note that MEA[Sulf]−MEA binary mixture is able to absorb CO2 from air at the room temperature. Also we found, that there were two temperature ranges of 323−423 K and 423−523 K corresponding to different binding character of carbon dioxide in the process of absorption/ desorption. The results of thermal desorption analysis were in agreement with CO2 solubility data, according to which the MEA[Sulf]− MEA binary mixture performed relatively high sorption capacity and three-stage desorption behavior pointing to the presence of two mechanisms of CO2 binding. MEA[Nic]− MEA also exhibited the relatively high CO2 solubility and twostage MEA desorption from two components of mixture. Therefore, summing up the experimental results, we conclude that the optimal absorbent for acid gases capture between the four investigated ones is the MEA[Sulf]−MEA binary mixture. The legible comparison indicates that the investigated binary mixtures have the absorption capacities close to the conventional ILs most widely studied for the capture of acidic gases. The highest values of the Henry’s Law constants for H2S, approximately equal to emim[Tf2N], bmim[Tf2N] and emim[eFAP] sorption capacity, were obtained for MEA[Sulf]−MEA and MEA[Nic]−MEA mixtures. The TEA-containing binary mixtures performed lower absorption properties comparable to hmim[Tf2N] and bmim[PF6]. In common, the conventional ILs spreading in industrial scales is limited by a high cost and toxicity of fluorinated molten salts. The approach to replace them by the protic ILs containing carboxylic anions could be competitive owing to the less harmful effects and economical profits combined with the similar sorption properties. The

effective physical sorption of gases in ILs occurs predominantly at elevated pressures, whereas the amine component introduction facilitates the gases capture at ambient conditions. Thus, the present work and absorbents described demonstrate one of the approaches to perform acidic gases capture by the hybrid solvents.

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CONCLUSIONS There H2S and CO2 solubility data in a series of alkanolamine− protic ionic liquid binary mixtures combining physical and chemical absorption mechanisms were presented. The binary mixtures were prepared by the neutralization reaction of ethanolamine or triethanolamine with 2-hydroxy-5-sulfobenzoic acid or pyridine-3-carboxylic acid, respectively. The increase of absorbent’s desorption temperature was expected owing to the effect of proton transfer from the cation. The density and the viscosity of four binary mixtures were compared to those of conventional ionic liquids. The viscosities of absorbents studied were relatively high, although the MEA-based absorbent performed the activation energy of viscous flow close to the value of those for bmim[PF6]. Two absorbents based on the monoethanolamine (MEA) component performed high absorption properties toward H2S and were able to bind acid gases through the dissolution of gas in protic ionic liquid and MEA component. The solubility of hydrogen sulfide, characterized by the Henry’s Law constant, in MEA-based binary mixtures was comparable to the commercially available ionic liquids. The results of thermal desorption analysis confirmed the capture of acid gases in MEA-based absorbent through the dissolution in MEA component and in protic ionic liquid. Although the alkanolamine−protic ionic liquid binary mixtures had relatively high densities and viscosities, the low cost and the high absorption capacity competing to the commercially available ionic liquids, will be attractive features for the further optimization of absorbents. In particular, the investigated absorbents can be tested for the selective removal 3435

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ACS Sustainable Chemistry & Engineering

(11) Lozano, P.; Bernal, J. M.; Sanchez-Gomez, G.; García-Verdugo, E.; Vaultier, M.; Burguete, M. I.; Luis, S. V. Sponge-like ionic liquids: a new platform for green biocatalytic chemical processes. Green Chem. 2015, 17, 3706−3717. (12) Tomé, L. C.; Marrucho, I. M. Ionic Liquid-Based Materials: A Platform to Design Engineered CO2 Separation Membranes. Chem. Soc. Rev. 2016, 45, 2785−2824. (13) Babamohammadi, S.; Shamiri, A.; Aroua, M. K. A Review of CO2 Capture by Absorption in Ionic Liquid-Based Solvents. Rev. Chem. Eng. 2015, 31, 383−412. (14) Fortunato, R.; Branco, L. C.; Afonso, C. A. M.; Benavente, J.; Crespo, J. G. Electrical Impedance Spectroscopy Characterisation of Supported Ionic Liquid Membranes. J. Membr. Sci. 2006, 270, 42−49. (15) Bazhenov, S.; Ramdin, M.; Volkov, A.; Volkov, V.; Vlugt, T. J. H.; de Loos, T. W. De. CO 2 Solubility in Biodegradable Hydroxylammonium-Based Ionic Liquids. J. Chem. Eng. Data 2014, 59, 702−708. (16) Otvagina, K. V.; Mochalova, A. E.; Sazanova, T. S.; Petukhov, A. N.; Moskvichev, A. A.; Vorotyntsev, A. V.; Afonso, C. A. M.; Vorotyntsev, I. V. Preparation and Characterization of Facilitated Transport Membranes Composed of Chitosan-Styrene and ChitosanAcrylonitrile Copolymers Modified by Methylimidazolium Based Ionic Liquids for CO2 Separation from CH4 and N2. Membranes (Basel, Switz.) 2016, 6, 31. (17) Akhmetshina, A. A.; Davletbaeva, I. M.; Grebenschikova, E. S.; Sazanova, T. S.; Petukhov, A. N.; Atlaskin, A. A.; Razov, E. N.; Zaripov, I. I.; Martins, C. F.; Neves, L. A.; et al. The Effect of Microporous Polymeric Support Modification on Surface and Gas Transport Properties of Supported Ionic Liquid Membranes. Membranes (Basel, Switz.) 2016, 6, 4. (18) Aziz, N.; Yusoff, R.; Aroua, M. K. Absorption of CO2 in Aqueous Mixtures of N-Methyldiethanolamine and Guanidinium Tris(pentafluoroethyl)trifluorophosphate Ionic Liquid at High-Pressure. Fluid Phase Equilib. 2012, 322−323, 120−125. (19) Greaves, T. L.; Drummond, C. J. Protic Ionic Liquids: Properties and Applications. Chem. Rev. 2008, 108, 206−237. (20) Yuan, X.; Zhang, S.; Liu, J.; Lu, X. Solubilities of CO2 in Hydroxyl Ammonium Ionic Liquids at Elevated Pressures. Fluid Phase Equilib. 2007, 257, 195−200. (21) Mattedi, S.; Carvalho, P. J.; Coutinho, J. A. P.; Alvarez, V. H.; Iglesias, M. High Pressure CO2 Solubility in N-Methyl-2-Hydroxyethylammonium Protic Ionic Liquids. J. Supercrit. Fluids 2011, 56, 224−230. (22) Yang, J.; Yu, X.; Yan, J.; Tu, S. T. CO2 Capture Using Amine Solution Mixed with Ionic Liquid. Ind. Eng. Chem. Res. 2014, 53, 2790−2799. (23) Yu, G.; Fan, S.; Chen, X.; Abdeltawab, A. A.; Al-Deyab, S. S. CO2 Absorption by Binary Mixture of Ionic Liquids-Monoethanolamine at Lower Pressure. Int. J. Greenhouse Gas Control 2016, 44, 52− 58. (24) Taib, M. M.; Murugesan, T. Solubilities of CO2 in Aqueous Solutions of Ionic Liquids (ILs) and Monoethanolamine (MEA) at Pressures from 100 to 1600 kPa. Chem. Eng. J. 2012, 181−182, 56−62. (25) Camper, D.; Bara, J. E.; Gin, D. L.; Noble, R. D. RoomTemperature Ionic Liquid-Amine Solutions: Tunable Solvents for Efficient and Reversible Capture of CO2. Ind. Eng. Chem. Res. 2008, 47, 8496−8498. (26) Ahmady, A.; Hashim, M. A.; Aroua, M. K. Absorption of Carbon Dioxide in the Aqueous Mixtures of Methyldiethanolamine with Three Types of Imidazolium-Based Ionic Liquids. Fluid Phase Equilib. 2011, 309, 76−82. (27) Osman, K.; Ramjugernath, D.; Coquelet, C. CO2 Solubility in Hybrid Solvents Containing 1-Butyl-3-Methylimidazolium Tetrafluoroborate and Mixtures of Alkanolamines. J. Chem. Eng. Data 2015, 60, 2380−2391. (28) Yu, Y.; Lu, X.; Zhou, Q.; Dong, K.; Yao, H.; Zhang, S. Biodegradable Naphthenic Acid Ionic Liquids: Synthesis, Characterization, and Quantitative Structure-Biodegradation Relationship. Chem. - Eur. J. 2008, 14, 11174−11182.

of hydrogen sulfide over carbon dioxide in the case of biogas purification and in the process of natural gas purification to improve the efficiency of sulfur recovery units.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acssuschemeng.7b00092. Results of NMR and CO2 and H2S solubility measurements at temperatures of 303.2, 313.2, 323.2, 333.2 K and pressures of up to 12.7 bar (PDF)

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AUTHOR INFORMATION

Corresponding Author

*Ilya V. Vorotyntsev, E-mail: [email protected]. ORCID

Ilya V. Vorotyntsev: 0000-0003-2282-0811 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the Russian Science Foundation (grant no. 15-19-10057). A.I.A. thanks PhosAgro/UNESCO/ IUPAC for their respective 2016 Green Chemistry for Life Grant. The authors are grateful to product manager Oksana Karabelskaya, service engineer Dmitry Levin from JSC “Aurora” for the assistance in the viscosity and density measurements.

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REFERENCES

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