Influence of Temperature on CO2 Absorption Rate and Capacity in

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Influence of Temperature on CO2 Absorption Rate and Capacity in Ionic Liquids M. P. Gimeno,* M. C. Mayoral, and J. M. Andrés Department of Chemical Processes, Instituto de Carboquímica (ICB-CSIC), Miguel Luesma Castán, 4, Zaragoza 50018, Spain ABSTRACT: Ionic liquids (ILs) present a new alternative for postcombustion CO2 capture because their properties can be tuned by means of different cation−anion combinations in order to obtain the most suitable properties for a specific task. This work is a novel study that investigates the absorption kinetics and evolution of the chemical transformations produced by CO2 absorption in ILs at different temperatures. A large range of pure ILs was tested with a screening process based on the pKa anion for efficient and reversible CO2 capture. CO2 uptake by selected ILs was determined for a wide range of pressures between atmospheric pressure and 2000 kPa at room temperature. Results show that CO2 absorption capacities of [bmim][Ac] and [bmim][Phen] at 100 kPa are close to those obtained at higher pressures, suggesting the existence of two solvation regimes. This was confirmed by IR analysis. The kinetics were significantly affected by the temperature and were shown to increase sharply with it. The optimum temperature for CO2 capture in [bmim][Ac] and [bmim][Phen] was found to be 323.15 and 348.15 K, respectively. kJ/mol of CO 2 at 313.15 K). 8,9 Therefore, although modifications have been made to the available amine-based technologies in order to offer better performance,2,10 new methods are required. To overcome these problems, several new technologies have been proposed as alternatives to the amine processes for CO2 capture from flue gases.2 Over the past decade, ionic liquids (IL) have garnered remarkable interest, both in academia and industry fields. It is based on the large versatility of these new fluids.4 Currently, ILs are presented as new alternative absorbents to capture CO2.11−13 Ionic liquids are organic salts with a melting point below 373.15 K and decomposition temperatures of around 573.15−673.15 K. ILs are an attractive option because their properties can be tuned by means of different cation−anion combinations in order to get the most suitable properties for a specific task.14 Anderson et al.15 suggest that there are about 1018 possible structures for room temperature ionic liquids. ILs have several advantages over other solvents, such as amines, that make them excellent for CO2 capture from gas mixtures: low melting point, almost negligible vapor pressure, wide temperature liquid range, viscosity higher than organic fluids but decreasing remarkably with temperature, low heat of reaction, low thermal and oxidative degradation, and tunable properties through an adequate combination of ions (task specific fluids), in addition to CO2 being highly soluble in ILs (much more than other common gases). To date, many research groups have studied CO2 solubility in ILs, and important research work is being done to improve IL/ CO2 systems.16−37 Bates et al.25 studied CO2 absorption by an IL based on a [BF4] cation and an imidazolium cation, which is functionalized with a primary amine. In this case, the CO2 absorption takes place through a reversible chemical process

1. INTRODUCTION There is growing consensus regarding the existence of worldwide climate change, and the anthropogenic emission of greenhouse gases (GHGs) into the atmosphere are considered as the major cause.1 It is widely accepted that one of the most important GHGs is carbon dioxide from energy production Therefore, most efforts are being directed toward reducing emissions generated by the largest sources, mainly power stations and industries that burn fossil fuels. Carbon capture and storage (CCS) technologies are emerging as a promising near-term path for reduction in CO2 emissions. Depending on the process or power-generating facility to which these technologies are applied, there are three main approaches to capturing the CO2 produced: precombustion, postcombustion, and oxyfuel combustion. Of these, postcombustion treatment is the most suitable option for retrofitting in existing fuel powered plants without major modifications.2,3 There are several available technologies for postcombustion CO2 capture. However, amine-based CO2 capture appears to be a viable option to make a significant reduction in these emissions.4 In these methods, the liquid absorbent is a mixture of water and an amine (primary, secondary or tertiary), usually monoethanolamine or diethanolamine. A chemical absorption of CO2 is produced, in which the amines used react with CO2 and lead to carbamate formation.5 This technology has been used for years to remove acid gases (H2S) from natural gas, and its use allows the removal CO2 from mixed gases at low partial pressure, as in flue gases.6 Nevertheless, the capture capacity of these methods is affected by different issues, as equilibrium limits and changes over the amine solution produced in the regeneration and recycling processes. Among these changes, it can be mentioned volatilization and degradation, leading to a number of corrosion and environmental problems with important increases in operational costs.7 Moreover, the primary drawback is thermodynamic limitation owing to the large amount of energy required to decompose the carbamate during regeneration (85 © XXXX American Chemical Society

Received: April 18, 2013

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where the CO2 is fixed as ammonium carbamate. CO2 absorption achieved by this IL was close to that obtained in amine processes. Anthony et al.20−26 measured CO2 solubility in ionic liquids 1-n-butyl-3-methylimidazolium hexafluorophosphate, [bmim][PF6], and 1-n-butyl-3-methylimidazolium tetrafluoroborate, [bmim][BF4], and compared their results with conventional technology using MEA. The authors noticed that although physical absorption took place, the Henry’s law constants were so high that the CO2 absorption through this way was no competitive with the traditional process based on CO2 capture with amines. A ternary equation of state (EOS) model was developed by Shiflett et al.27 for the N2O/CO2/ [bmim][BF4] system in order to understand separation of these gases using room-temperature ionic liquids (RTILs). Baltus et al.23 determined the CO2 absorption over an IL formed by an alkil imidazolium cation and a bis[trifluoromethylsulfonyl]amide anion, [Tf 2 N −]. These authors found that CO 2 absorption increased when the chain length of the imidazolium ring was increased. This effect was more accused in ionic liquids with [Tf2N−] than those with [PF6−] anions. On the other hand, they observed that imidazolium-based ionic liquid with a fluorine-substituted octyl side chain increase the CO 2 absorption. Pérez-Salado et al. and Jacquemin et al.24,35 measured CO2 solubility in [bmim][PF6] ionic liquid. The CO2 content in flue gas was impossible to be reduced to the low levels that were required. The main cause was that the Henry’s law constants were too large. Several authors have studied CO2 absorption by [bmim][Ac].29,31,34 Shiflett et al.29 compared the energy requirement and economic investment of a commercial MEA PC CO2 capture facility with a new process designed to use ionic liquids. Their results showed that, in both cases, more than 90% of CO2 could be removed and that the IL (1-butyl-3methylimidazolium acetate) process could reduce energy losses by 16%, compared to the MEA process. Besnard et al.31 studied the formation of 1-butyl-3methylimidazolium-2carboxylate in the mixture of CO2 with [bmim][Ac] at molecular level by means of NMR spectroscopy and DFT calculations. They found that the ions interact with the carboxylate molecule present in the IL structure. Their results showed that CO2 reacted with [bmim][Ac] to form 1-butyl3methylimidazolium-2carboxylate producing acetic acid, which was originated from the release of the proton bonded to carbon 2 of the imidazolium ring. The authors concluded that a bond between the CO2 and the imidazolium ring was produced. Carbaço et al.34 interpreted the interaction of CO2 solubility in [bmim][Ac], which resulted from two different solvation regimes. The first of these was controlled by an irreversible chemical reaction, ending in the imidazolium ring carboxylation and the acetic acid formation. The second corresponded to a physical absorption of the CO2. Gurau et al.30 reported experimental results about the absorption of CO2 by acetate ILs. It was found that the products obtained formed a singlecrystal X- ray structure. Mixtures of RTILs and alkanolamines were also studied.36 It was observed that CO2 capture in RTIL−amine solutions occurred rapidly and was readily reversed and that its behavior was similar to that found in aqueous amine solutions. However, more research is needed in this field. Wang et al.37 developed diverse phenolic ILs by neutralizing phosphonium hydroxide with various substituted phenols. The authors concluded that it was possible to achieve good absorption capacity as well as low absorption enthalpy by varying the substituents on the anion. They observed that physical properties (decomposition, viscosity, etc.) were

affected by the substituents on the phenolic ILs. Their results indicated that electron-withdrawing or electron-donating ability, position and quantity of substituents on the anion affect CO2 absorption. Several experimental methods based on pressure change measurement and volumetric and gravimetric analyses are currently being used in order to determine the CO2 absorption capabilities by ionic liquids. Of these, thermogravimetric analysis seems to be a fast, simple, and cheap technique for the accurate determination of CO2 absorption capacity and rate.26,33,38−42 The aim of this work is to compare the absorption capabilities of diverse ILs based on the 1-butyl-3-methylimidazolium cation [bmim+] and 14 anions. For this study, the CO2 absorption was determined for a wide range of pressures between 100 and 2000 kPa and at 303.15 K. The obtained results for the selected IL were also compared with a reference absorption process using a 50 wt % MEA solution as the solvent. This is the first time the evolution of CO2 loading has been measured in a study in order to determine CO2 absorption kinetics. For this purpose, the evolution of the CO2 absorption was precisely determined through an IR reaction cell top plate and by means of thermogravimetric analysis.

2. EXPERIMENTAL SECTION 2.1. Materials. The ILs used in this work were synthesized in our laboratory by metathesis in aqueous phase,43 with the exception of the 1-butyl-3-methylimidazolium tetrafluoroborate [bmim][BF4], which was purchased from Fluka Analytical (≥97.0%). The structure of the ILs used are shown in Figure 1. Chemicals and suppliers were as follows: sodium trifluoromethanesulfonate (98%), bis(trifluoromethane)sulfonamide lithium salt (≥99%), 1-butyl-3-methylimidazoium chloride (≥98%), 1-naphthoic acid (96%), and pentafluorophenol (≥99%) were obtained from Sigma-Aldrich. Tetrabutylammonium perfluoro-octanesulfonate (≥90%), methanesulfonic acid (≥99%), sodium dodecylsulfate (>97%), and phthalic acid (>99.5%) were purchased from Fluka Chemika. Silver oxide (extra pure), sodium hydroxide (≥99.9%), dichloromethane (99%), and monoethanolamine (≥98%) were supplied by Scharlau. Oxalic acid (99.5%) and phenol (98.5%) were obtained from Panreac. Before the study, the ionic liquids were degassed and dried under vacuum (P = 10 kPa and T = 373.15 K) over a period of about 24 h to remove traces of water and other volatile impurities. CO2 was supplied by Carburos Metálicos S.A. with a purity of ≥99.999%. 2.2. Apparatus and Procedure. The samples were characterized by IR spectroscopy, thermogravimetric analysis (TGA), and ion chromatography. The infrared spectra of ILs were obtained with an ABB Bomem MB3000 IR spectrometer between 600 cm−1 and 4000 cm−1 using a Pike’s Miracle attenuated total reflectance device (ATRIR) to gain a better understanding of the ILs formation and to detect the interaction between the IL and CO2. The thermal stability of ILs or the decomposition temperature measurements was determined with a TA Instruments SDT Q600 thermogravimetric analyzer. The samples (between 35 mg and 40 mg) were run in aluminum pans under argon atmosphere by heating from 303.15 to 873.15 K at a rate of 20 K/min and at a flow rate of 100 mL/min to determine volatility, boiling point, and thermal decomposition. The onset temperature is reported as the intersection of the baseline weight and the tangent of the weight versus temperature curve when the decomposition takes place. The IL samples were also characterized by ion chromatography (Metrohm IC) using a Metrosep A Supp5 of 250 × 4.0 mm column to determinate the chloride content and, hence, the purity of the ILs. CO2 uptake by selected ILs was determined for a wide range of pressures between atmospheric pressure (100 kPa) and 2000 kPa and at 303.15 K. CO2 capture at 100 kPa and 303.15 K was performed by bubbling CO2 through the IL at 100 kPa and 303.15 K while stirring at B

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Determination of CO2 absorption capability at high pressure was performed in a Thermo Cahn high pressure microbalance. One isothermal step was applied at 303.15 K, and the pressure range studied was from 500 to 2000 kPa. A dynamic mode operation was used in order to study the effect of the pressure; in this case, a continuous CO2 flow of 1500 mL/min at the desired pressure passed the sample while the weight increase was recorded. Solubility was meassured at a given pressure and temperature through the mass uptake (corrected for buoyancy effects) at each pressure. Temperature and pressure are precise to within ±0.5. The evolution of the chemical transformations produced by the CO2 absorption in the ILs at different temperatures was determined through the analysis of the infrared spectra obtained between 600 cm−1 and 4000 cm−1 from an attenuated total reflectance infrared (ATR-IR) BRUKER Tensor 27 using an autoclave-type reaction cell plate on top of a Golden Gate (Specac) ATR device. In this case, the IL was introduced into the reaction vessel and a continuous CO2 flow was passed through the cell for 24 h in order to avoid problems related with diffusion control. The temperature of the reaction cell top plate was controlled by use of a dedicated 4000 Series controller unit. Prior to the capture experiment, the IL was degassed at 423.15 K for 2 h and cooled to the required temperature.

3. RESULTS AND DISCUSSION 3.1. Characterization of ILs. The purity of the synthesized ILs was assessed by ion chromatography, the results of which are summarized in Table 1 together with their decomposition Table 1. Purity and Decomposition Temperature of the Synthesized Ionic Liquids ionic liquids

ILs (wt %)

Td (K)

[bmim][Cl] [bmim][NTFS] [bmim][TF] [bmim][PFOS] [bmim][DDBS] [bmim][MS] [bmim][Ox] (2:1) [bmim][Ox] (1:1) [bmim][Ac] [bmim][1-naf] [bmim][Phen] [bmim][PhenF5]

0.00 99.68 99.36 99.30 95.79 95.85 98.40 98.84 99.07 98.46 82.34 95.93

596.0 697.1 700.7 631.4 694.2 686.5 506.6 499.8 511.6 513.2 483.6 n/a

temperature, which was determined as the onset of the thermogravimetric curve. [Bmim][Ft] was previously excluded due to its extremely high viscosity. An important trend could be deduced from the experimental decomposition temperature data by taking into account the basicity of the anion. As seen, when the pKa value of the anion decreased, so did the decomposition temperature and, therefore, the IL stability. Further support to this hypothesis was provided by Wang et al.37 The stability of [bmim][PhenF5] could not be measured owing to analysis problems caused by a volume expansion. The nature of the anion also had a strong influence on the synthesis process and therefore on the IL purities that were achieved. IL purities presented a relationship with the pKa value and with the length of the fluoroalkyl chain on the anion, and it increased when the anion had a low pKa value ([1-naf−]) or a high length of the fluoroalkyl chain ([NTFS−]), which reduced the Lewis basicity of the carbonyl oxygen atoms caused by the electronegative fluorine atoms. 3.2. CO2 Capture. A series of 1-butyl-3-methylimidazoliumbased ILs in combination with diverse anions, with a wide range

Figure 1. Names, abbreviations, and ionic structures of ionic liquids used in this study.

moderate speed to enhance and maintain uniform solubility. A sample was withdrawn from the IL/CO2 at regular time intervals and analyzed thermogravimetrically (TA Instruments SDT Q600) to determine CO2 desorption and, therefore, the evolution of the CO2 absorption. CO2 absorption capabilities at different temperatures were also followed by static thermogravimetric analysis (TA Instruments SDT Q600). For this purpose, the samples were placed in aluminum pans under CO2 atmosphere by heating to the studied temperature (303.15, 323.15, or 348.15 K) at a rate of 278.15 K/min and at a flow rate of 100 mL/min. An isothermal step of 10 h was then applied to determine the maximum CO2 absorption capability of the IL. Prior to determining the CO2 absorption capabilities and, thus, to introducing the CO2, the ILs were heated to 423.15 K with N2 in order to removed possible volatile impurities. C

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of chemical types and basicities, were prepared by metathesis in aqueous phase. Study was made of the experimental solubilities of CO2 in pure ILs and pure MEA and an aqueous solution of MEA (50:50 vol. as should be used in industrial processes) at different pressures (between 100 and 2000 kPa) and at 303.15 K. The main purpose of this was to obtain information about the effect of different anions and operating conditions on CO2 absorption, as well as to obtain results of CO2 capture kinetics at different temperatures (from 303.15 to 373.15 K). The effect of changing the anion of the IL is significant, as can be observed in Figure 2, where the higher weight increases

Table 2. Absorption Capacity of CO2 in the Synthesized ILs at 2000 kPa, 303.15 K, and 60 min for all the Studied ILs ionic liquids

CO2 absorption (χCO2)

weight increase (wt %)

[bmim][NTFS] [bmim][TF] [bmim][PFOS] [bmim][DDBS] [bmim][MS] [bmim][Ox] (2:1) [bmim][Ox] (1:1) [bmim][Ac] [bmim][BF4] [Bmim][Ft] [bmim][1-naf] [bmim][Phen] [bmim][PhenF5] MEA MEA (50:50 vol)

0.30 0.22 0.07 0.14 0.08 0.08 0.03 0.38 0.13 n/a 0.08 0.27 n/a 0.01 0.30

4.69 4.24 0.51 1.59 1.53 0.93 0.54 13.31 2.86 n/a 1.17 6.87 n/a 0.87 11.98

[bmim][MS] > [bmim][PFOS] > [bmim][Ox] (1:1), while their CO2 absorption capacity was considerably slower than any of the other studied ILs. In order to study and compare the effect of absorption pressure, CO2 absorptions were measured at different pressures (2000, 1000, and 500 kPa) and 303.15 K. For this purpose, the ILs that showed the highest CO2 absorption capacities were selected ([bmim][NTFS], [bmim][Ac], [bmim][TF], and [bmim][Phen]). The results are shown in Table 3, and some Figure 2. Effect of the anion of ILs on the CO2 absorption capacity as a function of time by TGA at 2000 kPa and 303.15 K. Ionic Liquids: (1) MEA (50:50 vol.); (2) [bmim][Ac]; (3) [bmim][NTFS]; (4) [bmim][Phen]; (5) [bmim][BF4]; (6) [bmim][DDBS]; (7) [bmim][1-naf]; (8) MEA.

Table 3. CO2 Absorption Capacity in Various ILs at Several Pressures (2000, 1000, and 500 kPa), 303.15 K and 180 min CO2 absorption (χCO2)

obtained for the CO2 capture by the ILs are shown. As expected, the absorption of pure ILs was observed to be slower than that of aqueous solution of MEA at 60 min. This behavior could be explained by the fact that CO2 diffusivity and thus the rate of absorption is determined by the viscosity of the solvent, which is higher for the ILs than for the aqueous solution of MEA. This problem may be overcome by the addition of water since viscosity decreases when the water content is increased, speeding up the absorption rate, as can be seen by comparing the behavior of pure MEA with that of an aqueous solution of MEA. However, it is important to highlight that, while at 60 min the aqueous solution of MEA had reached phase equilibrium, indicated by the unchanging weight of the sample over time, the ILs had not yet reached their phase equilibrium as a result of their slower CO2 absorption reaction kinetics and, therefore, had a higher CO2 absorption capacity. The measured CO2 absorption capacity in the synthesized ILs at 303.15 K and 2000 kPa for all the studied ILs at 60 min are presented in Table 2. The absorption capacity results are expressed as mole fraction of CO2 in CO2−IL mixtures and as weight increase. Although the absorption kinetics of the studied ILs was slower than that of the MEA aqueous solution, the total amount of absorbed CO2 (Table 2, molar fraction) by [bmim][NTFS] and [bmim][Ac] at 60 min was similar to that of the MEA solution. The [bmim][TF] and [bmim][Phen] had the second highest absorption capacity, and [bmim][BF4] and [bmim][DDBS] show a slightly lower solubility. The remaining five ILs had a similar order: [bmim][Ox] (2:1) > [bmim][1-naf] >

ionic liquids

P = 2000 kPa

P = 1000 kPa

P = 500 kPa

[bmim][NTFS] [bmim][TF] [bmim][Ac] [bmim][Phen] MEA (50:50 vol)

0.30 0.22 0.38 0.27 0.39

0.14 0.13 0.36 0.19 0.35

0.07 0.06 0.32 0.17 0.32

of them are depicted in Figure 3. [bmim][NTFS] and [bmim][TF] presented a linearity with gas pressure, and as was expected, these ILs exhibited only physical solubility in accordance with Henry’s law. However, the increase in CO2 solubility for [bmim][Ac] and [bmim][Phen] were not proportional to the pressure increase, which indicates that a chemical reaction should take place in the absorption process. The initial absorption rate for these ILs was found to be higher when the pressure was increased (Figure 3) because no significant change was caused to the conformational alignment of the IL by the interactions between the IL anion and the CO2 molecules.16 It is important to notice that absorption pressure affected the CO2 absorption capacity of the [bmim][Ac] and [bmim][Phen] equally. However, the effect of absorption pressure over the initial absorption rate was more marked in the case of [bmim][Phen], indicating that the CO2 capture kinetics in [bmim][Phen] was slower than that occurring in [bmim][Ac]. CO2 absorption capacity was then studied at 100 kPa, 303.15 K, and 60 min in order to simplify the operating conditions in an industrial process (Table 4); [bmim][Ac] and [bmim][Phen] showed the highest CO2 absorption capacities at D

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chemical absorption. The other studied ILs seems to obey Henry’s law, with CO2 absorption capacities decreasing sharply with the pressure, corresponding to low physical absorption. For the anions that contain fluoroalkyl groups, it was found that viscosity increased with increasing amounts of F groups on the anion. This effect led to a sharp decrease in CO2 absorption solubility; that is, CO2 solubility in [bmim][NTFS] decreased from 0.300 at 2000 kPa to 0.030 at 100 kPa. The existence of chemical absorption was studied by comparison of the IR spectra of CO2−IL mixtures with those of pure IL. This chemical reaction is produced by the bond of CO2 with the imidazolium ring as result of the hydrogen substitution of its carbon 2. Figure 4 shows [bmim][Ac] and

Figure 3. Effect of absorption pressure on the CO2 absorption capacity by TGA at 2000, 1000, and 500 kPa and 303.15 K for [bmim][Ac] and [bmim][Phen].

Table 4. CO2 Absorption Capacity in Various ILs at 100 kPa, 303.15 K, and 60 min ionic liquids

CO2 absorption (χCO2)

[bmim][NTFS] [bmim][TF] [bmim][PFOS] [bmim][DDBS] [bmim][MS] [bmim][Ox] (1:1) [bmim][Ac] [bmim][BF4] [bmim][1-naf] [bmim][Phen] [bmim][PhenF5]

0.03 0.03 0.04 0.07 0.05 0.04 0.24 0.02 0.01 0.20 0.02

Figure 4. IR spectra of pure IL, CO2−IL mixtures, and the IR spectra resulting from the difference between pure IL and CO2−IL mixture. IL: (a) [bmim][Ac]; (b) [bmim][Phen].

atmospheric pressure, which were close to those obtained at high pressures (2000 kPa). This increased solubility has been assigned to the presence of two solvation regimes, one of them occurring through chemical interactions between acidic CO2 and the basic anion, and the other taking place through physisorption. On the other hand, [bmim][1-naf] showed the lowest CO2 absorption capacity. It is thought that this may be related to the pKa of the anion (3.7), which is close to the pKa of CO2 (3.58), hindering its displacement and preventing

[bmim][Phen] IR spectra from 600 cm−1 to 4000 cm−1, allowing characteristic absorption bands to be assigned to the presence of a carboxylate salt (COO−). In the 600−2000 cm−1 regions, three new peaks at 794 cm−1 (δ OCO), 1325 cm−1 (ν COO symmetric stretch), and 1665 cm−1 (ν COO asymmetric stretch) were identified with the group vibration of the functional COO group of the aromatic carboxylate molecule, which were in good agreement with different authors.29,31 Both E

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IR spectra showed a peak at 2336 cm−1, which corresponded to the vibration frequency of the CO2 molecule. This could be explained by the fact that a flow of CO2 was fed through the reaction cell during the IR spectra measurements and that it was not the result of a possible CO2 desorption. No differences were found between pure [bmim][1-naf] or [bmim][Ox](2:1) and its CO2−IL mixture, which was in good agreement with the experimental CO2 absorption capacity values. Based on the results of the screening experiments, we predominantly chose ILs that present chemical absorption ([bmim][Ac] and [bmim][Phen]) in order to obtain results on CO2 capture kinetics at different temperatures. To determine CO2 capture kinetics at different temperatures, the absorption band at 1666 cm−1 was followed; this corresponded to aromatic carboxylate salt formation. Figure 5a shows the evolution of the

achieved at 323.15 K regardless of the temperature studied, which only affected CO2 absorption capacity. The time required to achieve the thermodynamic equilibrium was longer for the [bmim][Ac] at 373.15 K. The experimental results obtained by TGA are presented in Table 5. It is interesting to Table 5. Amount of CO2 Absorbed in [bmin][Ac] and [bmin][Phen] at Different Temperatures and 600 min temp. (K)

[Bmin][Ac] (χCO2)

[Bmin][Phen] (χCO2)

303.15 323.15 348.15 373.15

0.27 0.20 0.14 0.04

0.20 0.04 0 0

notice that a CO2 absorption capacity deviation was observed with respect to the experimental values determined by TGA. In this case, the greatest CO2 absorption capacity was obtained at 303.15 K. This discrepancy could be attributed to small differences between the contact mode in the reaction cell and in TGA and the contribution of both solvation regimes, physical and chemical absorption. It could also have been caused by the acetic acid evaporation, which was produced during the carboxylation reaction. However, further investigation is needed. Figure 6a presents the evolution of the [bmim][Phen] absorbance over time for several temperatures. It can be seen that there was a significant change in CO2 absorption owing to an increase in temperature. From the screening experiments, it could be deduced that optimum CO2 absorption capture took place at 348.15 K. Faster absortion rates were observed when temperature was incresased, especially for temperatures of 348.15 and 373.15 K, and the highest CO2 absorption was achieved within first 30 and 60 min, respectively. The maximum CO2 absorption at 323.15 K was reached at 5 h, whereas equilibrium was not achieved for lower temperatures (303.15 K) even after 24 h. This could be assigned to the decrease in viscosity and to an increase in diffusivity at higher temperatures. According to the TGA results, the maximum CO2 absorption capacity in [bmim][Phen] was reached at 348.15 K (Figure 6b), which was corroborated by the absorbance results. However, TGA of [bmim][Phen] showed that weight was not constant throughout the entire experiment, especially when higher temperatures were used. This is in contrast with the absorbance bands obtained, which were fairly constant. Therefore, it could be thought to take place because the carboxylic acid displaces phenolate to phenol and then the phenol is evaporated form the mixture. The CO2 absorption capacity results obtained by TGA are shown in Table 5. Table 6 compares the CO2 absorption capacities obtained previously by other authors with those presented in this work. For this purpose, the ILs which showed the highest absorption capacities were selected. It can be seen that the CO2 absorption capacities obtained in this work are in the same order as those achieved by other authors. However, it is important to notice that the temperature used in the present study is, in most cases, lower than the others, which may have played a key role, given its influence over the IL viscosities and energy costs.

Figure 5. (a) Evolution of the [bmin][Ac] absorbance over time at different temperatures; (b) Evolution of the [bmin][Ac] weight increase over time at different temperatures.

[bmin][Ac] absorbance over time at different temperatures. In the studied range, a small increase in temperature of 25 K produced a strong increase in CO2 absorption capacity, while at higher temperatures the absorbance decreased sharply, owing to the limitation produced over the equilibrium conversion and decomposition by the thermodynamic equilibrium. The optimum temperature for CO2 capture in [bmim][Ac] was reached at 323.15 K, which is in good agreement with the results observed by thermogravimetric analysis (Figure 5b), where it can be seen that maximum CO2 absorption was

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CONCLUSION CO2 absorption capacity in a wide range of pure ILs was studied using a screening process based on the pKa anion. The results indicate that IL properties, such as decomposition F

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close to those obtained al high pressures. This was attributed to the existence of two absorption regimes, one of them controlled by chemical reaction and favored by the difference between the anion and CO2 pKa. New experimental data were reported for the absorption kinetics and the evolution of the chemical transformations produced by CO2 absorption in [bmim][Ac] and [bmim][Phen] in the temperature range between 303.15 and 373.15 K and at atmospheric pressure. In general, the absorption rate increased with an increase in temperature. When comparison was made between CO2 absorption capabilities at different temperatures, the results revealed that maximum CO 2 absorption was achieved at 303.15 and 348.15 K for [bmim][Ac] and [bmim][Phen], respectively. In an effort to gain a deeper insight into the chemical transformation produced by the CO2 absorption at different temperatures, IR spectra, determined as the difference between the pure IL and the CO2−IL mixture for [bmim][Ac] and [bmim][Phen], was measured. These results supported the chemical absorption of CO2 in the ILs by the assignation of the characteristic bands to the presence of the functional COO group of the carboxylate molecule.

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AUTHOR INFORMATION

Corresponding Author

*Tel: (+34) 976 733 977. Fax: (+34) 976 733 318. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS M.P. Gimeno acknowledges the support of the Spanish National Research Council (CSIC) through their JAE-Doc program.

Figure 6. (a) Evolution of the [bmin][Phen] absorbance over time at different temperatures; (b) evolution of the [bmin][Ac] weight increase over time at different temperatures.

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temperature, IL stability, and IL purity, were significantly affected by the pKa anion and increased when the pKa value of the anion decreased. The effect of absorption pressure (between 100 and 2000 kPa) at 303.15 K was also determined. For [bmim][Ac] and [bmim][Phen], the increase in CO2 solubility was not proportional to the pressure increase; therefore, a chemical reaction should be involved in the absorption process. Another issue was the solubility of CO2 at mild conditions (100 kPa and 303.15 K), where [bmim][Ac] and [bmim][Phen] showed higher CO2 absorption capacities

REFERENCES

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Table 6. CO2 Absorption Capacities for Different ILs ref Muldoon, M. J. et al.

33

Carbasso, M. I. et al.34

Chen, Y. et al.38 Gonzalez-Miquel, M. et al39

present work

T (K)

P (kPa)

[bmim][NTFS]

333.15 333.15 313.15 313.15 313.15 323.15 298.15 298.15 298.15 303.15 303.15 303.15

2000 1000 2000 1000 100 100 2000 1000 100 2000 1000 100

0.25 0.10

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[bmim][TF]

[bmim][Ac]

0.02

0.32 0.27 0.17 0.09

0.22 0.13 0.03

0.38 0.36 0.24

dx.doi.org/10.1021/ef401063r | Energy Fuels XXXX, XXX, XXX−XXX

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