Effect of Temperature and Water Concentration on CO2 Absorption by

Article pubs.acs.org/jced

Effect of Temperature and Water Concentration on CO2 Absorption by Tetrabutylphosphonium Formate Ionic Liquid Published as part of The Journal of Chemical and Engineering Data special issue “Proceedings of the 6th International Congress on Ionic Liquids” Yoshiro Yasaka* and Yoshifumi Kimura* Department of Molecular Chemistry and Biochemistry, Faculty of Science and Engineering, Doshisha University, Kyoto, Japan 610-0321 S Supporting Information *

ABSTRACT: The CO2 absorption capacities for binary solutions of tetrabutylphosphonium formate ionic liquid (IL) and water are measured at 0.1 MPa of CO2 for various compositions (up to 10 in mole ratio of water to the IL) as a function of temperature (−24 to 60 °C). The capacities measured as the mole of CO2 with respect to 1 mol of the IL varied in a wide range from 0.01 to 1.0. Capacities decreased monotonously with temperature at a fixed absorbent composition. When the water concentration is varied at a fixed temperature, capacity takes a maximum when the water mole ratio is approximately 1. In contrast to CO2 absorbents made from acetate ILs, the present system loses affinity to CO2 in the absence of water. The equilibrium constant for the chemisorption is defined in two ways by assuming that CO2 is captured as a free bicarbonate ion (HCO3−) or a complex ion with formic acid ([(HCOO) (HCO3)H] −) . The van’t Hoff plot of these two equilibrium constants are both linear. However, the estimated enthalpy change of absorption is ca. 20 kJ/mol larger for the former scenario than the latter.

1. INTRODUCTION Chemical absorption (chemisorption) of CO2 by ILs derived from carboxylate anions is attracting growing interest.1−16 The carboxylate ILs are unique CO2 chemisorbents in that seemingly they do not own strongly basic functional groups in contrast to popular CO2 chemisorbents such as amines17−19 and amino-functionalized ILs.20−26 Recently we reported CO2 absorption by tetrabutylphosphonium formate ([P4444][HCOO]) in the presence of water.2 It is remarkable that the formate anion is one of the least basic carboxylate anions but can capture CO2 even at 0.1 MPa. The [P4444][HCOO]−water system distinguishes itself among the other IL-based chemisorbents in that the absorption capacity drastically changes with the water content; little chemisorption is known at dry conditions3 but the capacity increases to more than 0.5 mol of CO2 per 1 mol of the IL when an equimolar water coexists.2 Taking advantage of this unique feature, we have proposed an absorption−desorption cycle based on the water-content sawing of the absorbent, which can be energetically favorable against the standard temperature sawing. From the engineering point of view, more data on absorption capacity as well as physical characteristics of the absorbent are indispensable for process optimization. The mechanisms of CO2 absorption by carboxylate ILs including formate ILs are still under discussion. Within the family of carboxylate ILs, 1-butyl-3-methylimidazolium acetate ([bmim][OAc]) can be the most studied one.4−9 However, the © XXXX American Chemical Society

chemisorption mechanism is rather specific to the chemical nature of the imidazolium cation (i.e., carbene formation) according to mechanistic studies.4,5 Carboxylate anions coupled with inert ammonium and phosphonium ions also absorb CO2 in the presence and/or absence of water but have been mostly dismissed. Pioneering work was done by Quinn and co-workers before “ionic liquids” came to widely acknowledged on several hydrates (salt/water = 1:1 to 1:4) of tetraethylammonium acetate and tetramethylammonium n-propionate.10,11 Their spectroscopic data were consistent with absorption of CO2 as bicarbonate. Closely related systems (triethylbutylammonium acetate/water = 1:1 to 1:5 and N-methyl-N-butylpyrrolidinium acetate/water = 1:0 to 1:4) have been studied by Wang et al.12,13 and Stevanovic et al.14 In our previous work,2 we measured the water content dependence of the absorption capacities of [P4444][HCOO]− water mixtures and elucidated that CO2 is chemisorbed by 1:1 stoichiometry with respect to water via the following reactions. CO2 + H 2O ⇌ H 2CO3

(1)

H 2CO3 + HCOO− ⇌ HCO3− + HCOOH

(2)

Received: August 14, 2015 Accepted: December 29, 2015

A

DOI: 10.1021/acs.jced.5b00694 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

Article

Table 1. Samples chemical name c

[P4444]Br NaBF4 formic acid cesium carbonate [P4444][HCOO]e CO2

grade

reported purity wt %a

method

NMR purity area %b

99.0 wt % 97 wt % 98 wt % 99.9 wt %

100.1 > 97 99.5 > 99.9

titration (AgNO3) precipitation titration (NaOH) trace metal analysis

99.98(31P) 99.43(19F) 99.998(1H)d

source Tokyo Kasei Nacalai Tesque Nacalai Tesque Aldrich Cabot high-purity grade synthesized Showa Denko Gas Products

> > > >

99.990 (31P)f > 99.990 vol %

a

Reported by the manufacturer on the certificate of analysis. bAnalyzed by the authors at the time of the receipt of the reagent.. cTetra-nbutylphosphonium bromide. dWater is not counted as impurity since it will be evaporated. eTetra-n-butylphosphonium formate. f0.02 mol % of fluorine impurity (in the form of BF4− and F−) was detected by 19F NMR. No impurity ( 1. When we examined the capacity in the water-poor region more closely (Figure 4c), it was found that it increases linearly to qH2O. The observation can be best interpreted by the 1:1 stoichiometry between CO2 and H2O. The results can be compared with the CO2 capacities of two tetraalkylammonium acetates reported by Quinn et al.11 and Wang et al.12 Interestingly, [N2224][OAc] can absorb as much CO2 in the absence of water as in its presence. This features a striking contrast in the water-concentration dependence of capacities between the formate and acetate ILs; the former shows capacity maxima in the mid composition, whereas the latter does not. However, the capacity of the [N2224][OAc]− water system decreases in a similar manor as that of formate system as the water−mole ratio exceeds 1. Taking advantage of the variation of absorption capacities with the water contents, we can operate the CO2 absorption/ desorption cycle by the addition/removal of water to/from the absorbent instead of changing the temperature. The energy cost of this process is dependent in one part on how sensitive the capacity is on the water content and in the other part how easy it is to remove water from the absorbent (water vapor pressure (activity) of the absorbent). One may be concerned that the absorbent may not be perfectly recycled when CO2 is desorbed at higher temperatures and/or reduced pressures if the formic acid formed in eq 2 in the Introduction evaporates. We addressed this issue by running the absorption−desorption cycle three times. The absorption was carried out at −10 °C and 0.1 MPa for 1 h. Desorption was carried out at 35 °C and 2.0 kPa for 0.5 h. From 13C NMR analysis, CO2 was loaded by a mole ratio of 0.4 at the absorption. The desorption allowed removal of ∼85% of CO2 that had been loaded. After each absorption−desorption cycle, 1H NMR was measured to check whether the absorbent is intact. The results are shown in Table 4. As can be seen, the mole ratio of formic acid plus the formate anion (NMR cannot distinguish these two) to the cation is perfectly preserved, whereas water is partly evaporated from the absorbent. Thus, the absorbent is reusable as long as water content is adjusted periodically. To focus on the efficiency of water-sawing cycle, absorption capacities are shown as normalized values in Figure 4b. The longitudinal axis represents the proportion of CO2 that remains in the absorbent after water is added to enlarge qH2O from 1 to an arbitral value. For the [P4444][HCOO]−water system, nearly 50% of CO2 is released when qH2O reaches 2−3 and nearly 80% of CO2 is released when qH2O reaches 4−6. The relative capacity more sharply decreases with increasing qH2O at 25 and 50 °C than at 0 °C. Compared to related acetate systems ([N2224][OAc]−water and [N2222][OAc]−water), the capacity dependence on qH2O of the [P4444][HCOO]−water system is regarded as typical for carboxylate ILs. 3.4. Comparison with Known CO2 Absorbents. In Figure 5 the CO2 absorption capacities at 0.1 MPa of CO2 are shown on a weight basis; the longitudinal axis is for the equilibrium loading of CO2 in grams per 1 g of each absorbent. Out of five carboxylate IL absorbents (with or without water)

a

Atmospheric pressure varied within the range of 985 to 1025 hPa due to meteorological reasons. The pressure at the time of each measurement (u(P) = 2 hPa) was obtained from the observation report by Kyoto Local Meteorological Office with −4 hPa correction for the sea level difference. u(T) = 0.1 K for T > 0 °C and 0.2 K for T < 0 °C. u(qH2O) = 0.02 for qH2O ≤ 0.6, and ur(qH2O) = 0.01 for qH2O ≥ 1. u(qCO2) = 0.02 for qH2O < 3, 0.01 for qH2O = 3.77 and 0.006 for qH2O ≥ 6. However, u(qCO2) is larger for T < 5 °C and T > 45 °C by a factor of 2. bMole ratio to the IL defined as qH2O = nH2O/nIL where ni is the mole of species i present in the absorbent. For nIL, the cation and anion are counted as a single component. E



Article

Figure 4. Water concentration dependence of the CO2 absorption capacity of [P4444][HCOO]−water mixtures. (a) Absorption capacity (qCO2) under CO2 of 0.1 MPa is plotted as the mole of CO2 divided by the mole of the IL contained in the absorbent. The water content (horizontal axis) is given as the mole of water divided by the mole of the IL contained in the absorbent. Data are shown for 0 (blue), 25 (green) and 50 °C (red). (b) Absorption capacity is normalized at qH2O = 1 for each temperature. Values for [N2222][OAc] and [N2224][OAc] are taken from refs 11 and 12, respectively. One data point (25 °C and in the absence of water) in [P4444][HCOO] is taken from ref 3. (c) The dependence of qCO2 ) on qH2O is shown for IL-rich compositions at 20 °C under CO2 of 0.1 MPa.

Table 4. 1H NMR Analysis after Repeated CO2 Absorption− Desorption Cycles on the [P4444][HCOO]−Water Absorbent no. of cycles

qH2Oa

qformb

initial 1 2 3

1.94 1.72 1.60 1.46

0.998 0.995 1.001 0.997

a Mole ratio of water measured as the peak integral for H2O divided by that for the tetra-n-butylphosphosnium cation. u(qH2O) = 0.02. bMole ratio of formic acid plus the formate anion measured as the peak integral for HCOOH/HCOO− (coalesced) divided by that for the tetra-n-butylphosphosnium cation. u(qform) = 0.005.

Figure 5. Comparison of the CO2 absorption capacities of [P4444][HCOO]−water mixtures (colored curves) with those of known CO2 absorbents on the weight basis. Capacities for dry [P4444][HCOO], [bmim][OAc] and [emim][OAc] are taken from ref 3 and those for [N2222][OAc] and [N2224][OAc] are taken from refs 11 and 12, respectively. Capacities for [bmim][OAc] in the presence of water (14 wt % or 1.5 equiv to the IL) are taken from ref 6. Capacity for the 30 wt % aqueous monoethanolamine (MEA) solution is measured in this work.

shown in the figure, the [P4444][HCOO]−water system exhibits slightly lower performance on a weight basis. This is mainly due to its relatively large molecular weight ([P4444][HCOO] = 304, whereas [N2224][OAc] = 217 and [bmim][OAc] = 198). The 30 wt % aqueous monoethanolamine solution (MEA) is often used as a benchmark for newly developed CO2 absorbents. We measured CO2 capacity of MEA in a temperature range of 5− 45 °C. Compared to MEA, the formate system exhibits lower capacities at room temperature by 50%. As can be seen from Figure 5, however, the MEA’s advantage is less pronounced at lower temperatures as the performance of the [P4444][HCOO]−water steeply improves.

4. DISCUSSIONS 4.1. Stoichiometry and Thermodynamics. We consider that the following points are especially informative from the F



Article

Figure 6. Van’t Hoff plots for the CO2 chemisorption equilibrium constants defined in (a) eq 9 and (b) eq 13. The water content (qH2O) of the absorbent is shown in the figure. The black star symbol in panel a indicates the equilibrium constant from aqueous chemistry.

it is obtained from Table 2. The variation of C0 due to a temperature change and a CO2 loading is not considered here. Alternately, we defined another equilibrium constant (K′) based on the reaction 6 by the following equation:

capacity data presented above for the [P4444][HCOO]−water system. (i). The capacities in terms of molar ratio are mostly bound above by the molar ratio of water. (ii). At low enough temperatures, the actual capacities are far beyond the stoichiometric capacity of 0.5 within the reaction mechanism involving the diformate anion formation (eqs 3 and 5 in the Introduction). In view of these the stoichiometry of chemisorption is most reasonably represented by the following equation: −

CO2 (g) + H 2O + HCOO ⇌

HCO3−

+ HCOOH

K′ =

where [HCOO(HCO3)H ] is evaluated by [HCOO(HCO3)H−] = C0 × qCO

2

(8)

A question arises here which mechanism better describes the equilibrium CO2 capacity data. To address this issue, we define equilibrium constants in two ways based on eqs 7 and 8 and analyze their temperature on a van’t Hoff plot. We define an equilibrium constant based on reaction 7 by the following equation:

ΔH = −R

(9)

where brackets indicate molar concentrations (activity coefficients are omitted) and PCO2 is fixed at 0.10 MPa in this study. In evaluating the K, we calculated the concentrations of relevant species from the measured capacity qCO2 by the following formulas: [HCOOH] = [HCO3−] = C0 × qCO

(10)

[HCOO−] = C0 × (1 − qCO )

(11)

[H 2O] = C0 × (qH O − qCO )

(12)

2

2

2

2

(14)

The obtained K and K′ values are plotted in Figure 6 for each absorbent composition as a function of temperature in the form of van’t Hoff plot. As is shown in the figures, the equilibrium constants for different water concentrations in the absorbent are nearly parallel and do not merge. Further the equilibrium constants are larger for water-poor (IL-rich) conditions. For a thermodynamically well-defined equilibrium constant, it is solely a function of temperature. In the case of K and K′, however, they are functions of water concentrations as well because we neglected the activity coefficients. It is a hard task to evaluate activity coefficeints in ILs, but we may speculate that variations of activity coefficients within the temperature range studied are small. Then we may roughly estimate heats of absorption from the slope of the plot; that is,

(7)

CO2 (g) + H 2O + HCOO− ⇌ [(HCOO)(HCO3)H]−

[HCOOH][HCO−3 ] [HCOO−][H 2O]PCO2

(13)

−

Further we may assume strong hydrogen-bond interactions between bicarbonate ion and formic acid (eq 4). In such a case, eq 7 is combined with eq 4 in the Introduction to give

K=

[HCOO(HCO3)H−] [HCOO−][H 2O]PCO2

∂(ln K ) ∂(1/T )

(15)

Interestingly the obtained values are quite different between the two definitions of the equilibrium constants (see Table 5). The ΔH values obtained from K is larger than that obtained from K′ by ca. 20 kJ mol−1. Therefore, we may check the formation of the complex by measuring the heat of absorption. This is a topic of future study. In any case, ΔH is larger for water-poorer compositions than water-rich compositions. The especially large value of ΔH (>65 kJ/mol) for qH2O < 1 suggests that basicity of formate ion and/ or bicarbonate anion are enhanced when it is less solvated (coordinated) by water. At the presence of limited moles of water, the solvation shell of the anions is only partially occupied by water. Thus, the negative charges are more localized than in aqueous solutions. This can facilitate the proton transfer from

−

where C0 = [HCOO ] + [HCOOH]. Since C0 is equal to the initial molar concentration of the formate ion in the absorbent, G



Article

an ideally behaved aqueous solution. The large CO2 absorption capacity certainly comes from a unique solvation environment for ions and molecules in ILs. 4.2. Effect of Anion Basicity. CO2 absorptions by acetate ILs have been reported in the literature. On the basis of pKa data from aqueous chemistry,27 acetate anion is by an order of magnitude more basic than formate anion. Therefore, one expects acetate ILs to possess higher affinity to CO2. Actually a striking anion effect is noticed in the absence of water. The CO2 capacities for neat acetate ILs are as follow: 0.47 in mole raio (0.06 MPa, 25 °C)12 for [N2224][OAc], 0.36 (0.1 MPa, 25 °C)3 for [emim][OAc], 0.37 (0.1 MPa, 25 °C)3 for [bmim][OAc], 0.27 (interpolated to 0.1 MPa, 30 °C)15 for [P4444][OAc]. In contrast, neat [P4444][HCOO] exhibited a capacity of 0.037 (0.1 MPa, 25 °C).3 N-Methyl-N-butylpyrrolidinium acetate showed exceptionally low (0.001 at 0.09 MPa, 80 °C)14 capacity as an acetate presumably due to the high temperature. These data are indicative of an active absorption mechanism apart from the bicarbonate mechanism shown in eq 2. Mechanistic studies have been carried out for some of the systems. The CO2 adduct of [bmim][OAc] was isolated and analyzed by X-ray crystallography.5 It was identified as an imidazolium derivative where CO2 is inserted in the C−H bond at the 2 position of the imidazole ring. Tetraalkylammonium and phoshphonium cations are considered as chemically inert. The large CO2 absorption capacity of [N2224][OAc] is reasoned as Lewis acid (CO2) and base (acetate anion) reactions;12 no spectroscopic measurements have been reported. For [P4444][OAc], a theoretical approach was practiced by Shi et al.15 by developing force fields that can reproduce attractive interactions between bent or straight CO2 molecules and acetate anions. A reasonable agreement was attained with the experimental CO2 capacities in their classical simulation. In the presence of water, spectroscopic investigations suggest that the bicarbonate mechanism (eq 2) is mainly operating for the acetate ILs as well. To be more specific, 13C NMR spectra reported for an equimolar mixture of [N2224][OAc] and water after CO2 absorption12 is quite similar to that observed by the present authors for the [P4444][HCOO]−water system.2 Upon the addition of water, the CO2−acetate interaction which dominates in the absence of water is presumably hampered by the coordination of water to the anion. It is not clear from the literature data at which water contents switching between these two absorption mechanisms may occur in acetate ILs. Absorption capacities are compared between the formate and acetate ILs in Figure 4a. [N2222][OAc] absorbs more CO2 than [P4444][HCOO] at 50 °C as is guessed from the higher basicity of the former. The [N2224][OAc], however, absorbs less CO2 when qH2O = 2. Apparently basicity of the anion is affected by the alkyl-chain length of the surrounding cations. It is to be noted that the performance of the aqueous ternary amine absorbents are directly correlated to the basicity of the amine.28 In view of this, the uniqueness of the IL absorbents originates from the richness of nonqaueous chemistry.

Table 5. Enthalpy of CO2 Absorption for [P4444][HCOO]− water Mixtures Obtained from the van’t Hoff Plot of the Chemisorption Equilibrium Constants K and K′ qH2Oa

ΔH/kJ mol−1 for Kb

ΔH/kJ mol−1 for K′c

75 84 86 61 53 58 54 58 49 48

55 66 65 46 41 34 32 32 26 24

d

0.40 0.50d 0.60e 1.0f 1.9f 2.3f 2.8f 3.9f 6.0f 9.9d a

Molar ratio to the IL defined as qH2O = nH2O/nIL. u(qH2O) = 0.02 for qH2O ≤ 0.6, and ur(qH2O) = 0.01 for qH2O ≥ 1 bK is defined as

K=

[HCOOH][HCO− 3] , [HCOO−][H2O]PCO2

where no complexation of formic acid is [HCOO(HCO3)H−] , where complexation [HCOO−][H2O]PCO2 d ion is assumed. ur(K) = 0.1 and ur(K′) f −1

c

assumed. K′ is defined as K ′ =

of formic acid and bicarbonate = 0.1. eur(K) = 0.03 and ur(K′) = 0.03. u(K) = 1 kJ mol 1 kJ mol−1.

and u(K′) =

carbonic acid to formate anion (eq 2) as well as complex formation in eq 4. For comparison, one may calculate the equilibrium constant defined in eq 9 based on aqueous chemistry. Literature data27 at 298 K yields K= =

[HCOOH][HCO−3 ] [HCOO−][H 2O]PCO2 K a,1(H 2CO3) 1 1 ·Khyd · · γ(CO2 ) K a(HCOOH) 55.5 mol/L

= 1.7 × 10−5 MPa−1

or

ln(K /MPa) = −10.96 where γ(CO2 ) =

PCO2 [CO2 (aq)]

= 2.6 MPamol−1L (physical dissolution)

Khyd =

[H 2CO3] 1 = (hydration of molecular CO2 ) [CO2 (aq)] 660

K a,1(H 2CO3) =

[HCO−3 ][H+] [H 2CO3]

= 2.8 × 10−4 mol L−1 (deprotonation for carbonic acid) K a(HCOOH) =

[HCOO−][H+] [HCOOH]

5. CONCLUSIONS Tetrabutylphosphonium formate exhibited large and reversible CO2 capture in the presence of water. The absorbed CO2 reacts with water to form the bicarbonate ion as deduced from the present capacity data. Unlike acetate ILs, the formate IL showed little CO2 absorption. This suggests that molecular CO2 has less affinity to formate ion than to acetate ion. Large

= 1.8 × 10−4 mol L−1 (deprotonation for formic acid)

When this aqueous equilibrium constant was plotted in Figure 6, it fell far below the plot for qH2O = 9.9. It is suggested that even at this compostion, the absorbent should not be treated as H



Article

(14) Stevanovic, S.; Podgorsek, A.; Moura, L.; Santini, C. C.; Padua, A. A. H.; Costa Gomes, M. F. Absorption of carbon dioxide by ionic liquids with carboxylate anions. Int. J. Greenhouse Gas Control 2013, 17, 78−88. (15) Shi, W.; Thompson, R. L.; Albenze, E.; Steckel, J. A.; Nulwala, H. B.; Luebke, D. R. Contribution of the acetate anion to CO2 solubility in ionic liquids: Theoretical method development and experimental study. J. Phys. Chem. B 2014, 118, 7383−7394. (16) Shiflett, M. B.; Yokozeki, A. Phase behavior of carbon dioxide in ionic liquids: [emim][Acetate], [emim][Trifluoroacetate], and [emim][Acetate] + [emim][Trifluoroacetate] mixtures. J. Chem. Eng. Data 2009, 54, 108−114. (17) Puxty, G.; Rowland, R.; Allport, A.; Yang, Q.; Bown, M.; Burns, R.; Maeder, M.; Attalla, M. Carbon dioxide postcombustion capture: A novel screening study of the carbon dioxide absorption performance of 76 amines. Environ. Sci. Technol. 2009, 43, 6427−6433. (18) Yang, H.; Xu, Z.; Fan, M.; Gupta, R.; Slimane, R. B.; Bland, A. E.; Wright, I. Progress in carbon dioxide separation and capture: A review. J. Environ. Sci. 2008, 20, 14−27. (19) 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. (20) Bates, E. D.; Mayton, R. D.; Ntai, I.; Davis, J. H. J. CO2 capture by a task-specific ionic liquid. J. Am. Chem. Soc. 2002, 124, 926−927. (21) Wang, C.; Luo, H.; Jiang, D.; Li, H.; Dai, S. Carbon Dioxide Capture by Superbase-Derived Protic Ionic Liquids. Angew. Chem., Int. Ed. 2010, 49, 5978−5981. (22) Zhang, Y.; Zhang, S.; Lu, X.; Zhou, Q.; Fan, W.; Zhang, X. Dual amino functionalized phosphonium ionic liquids for CO2 capture. Chem. - Eur. J. 2009, 15, 3003−3011. (23) Gurkan, B. E.; de la Fuente, J.; Mindrup, E. M.; Ficke, L. E.; Goodrich, B. F.; Price, E. A.; Schneider, W. F.; Brennecke, J. F. Equimolar CO2 absorption by anion-functionalized ionic liquids. J. Am. Chem. Soc. 2010, 132, 2116−2117. (24) Kasahara, S.; Kamio, E.; Ishigami, T.; Matsuyama, H. Amino acid ionic liquid-based facilitated transport membranes for CO2 separation. Chem. Commun. 2012, 48, 6903−6905. (25) Kasahara, S.; Kamio, E.; Matsuyama, H. Improvements in the CO2 permeation selectivities of amino acid ionic liquid-based fascilitated transport membranes by controlling their gas absorption properties. J. Membr. Sci. 2014, 454, 155−162. (26) Sanchez, L.; Meindersma, G.; De Haan, A. Kinetics of absorption of CO2 in amino-functionalized ionic liquis. Chem. Eng. J. 2010, 166, 11−1115. (27) CRC Handbook of Chemistry and Physics, 59th ed.; Weast, R. C., Ed.; CRC Press: West Palm Beach, FL, 1978−1979. (28) Tomizaki, K.; Shimizu, S.; Onoda, M.; Fujioka, Y. Heats of reaction and vapor−liquid equilibria of novel chemical absorbents for absorption/recovery of pressurized carbon dioxide in integrated coal gasification combined cycle−carbon capture and storage process. Ind. Eng. Chem. Res. 2010, 49, 1214−1221.

variations of absorption capacities with temperature and water concentrations render the formate ILs as a candidate for lowcost CO2 sequestration process.

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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/acs.jced.5b00694. NMR characterizations of synthesized materials, schematics of the apparatus, method of volumetric analysis, and error estimation (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. *E-mail: [email protected]. Funding

This work was supported by ENEOS Hydrogen Trust Fund and grants from Doshisha University. Notes

The authors declare no competing financial interest.

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REFERENCES

(1) Huang, J.; Rüther, T. Why ionic liquids attractive for CO2 absorption? An overview. Aust. J. Chem. 2009, 62, 298−308. (2) Yasaka, Y.; Ueno, M.; Kimura, Y. Chemisorption of carbon dioxide in carboxylate-functionalized ionic liquids: A mechanistic study. Chem. Lett. 2014, 43, 626−628. (3) Yokozeki, A.; Shiflett, M. B.; Junk, C. P.; Grieco, L. M.; Foo, T. Physical and chemical absorptions of carbon dioxide in roomtemperature ionic liquids. J. Phys. Chem. B 2008, 112, 16654−16663. (4) Maginn, E. G. Quarterly Technical Report to U. S. DOE, 01/05− 03/05; DOE Scientific and Technical Information: Oak Ridge, TN, 2005; pp 1−12. (5) Gurau, G.; Rodriguez, H.; Kelley, S. P.; Janiczek, P.; Kalb, R.; Rogers, R. D. Demonstration of chemisorption of carbon dioxide in 1,3-dialkylimidazolium acetate ionic liquids. Angew. Chem., Int. Ed. 2011, 50, 12024−12026. (6) Chinn, D.; Vu, D.; Driver, M. S.; Boudreau, L. C. CO2 removal from gas using ionic liquid absorbents. U.S. Patent 2005/0129598, 2005. (7) Stevanovic, S.; Podgoršek, A.; Padua, A. A. H.; Costa Gomes, M. F. Effect of water on the carbon dioxide absorption by 1-alkyl-3methylimidazolium acetate ionic liquids. J. Phys. Chem. B 2012, 116, 14416−14425. (8) Cabaço, M. I.; Besnard, M.; Danten, Y.; Coutinho, J. A. P. Carbon Dioxide in 1-Butyl-3-methylimidazolium Acetate. I. Unusual Solubility Investigated by Raman Spectroscopy and DFT Calculations. J. Phys. Chem. A 2012, 116, 1605−1620. (9) Shiflett, M. B.; Drew, D. W.; Cantini, R. A.; Yokozeki, A. Carbon Dioxide Capture Using Ionic Liquid 1-Butyl-3-methylimidazolium Acetate. Energy Fuels 2010, 24, 5781−5789. (10) Quinn, R.; Appleby, J. B.; Pez, G. P. Salt hydrates: New reversible absorbents for carbon dioxide. J. Am. Chem. Soc. 1995, 117, 329−335. (11) Quinn, R. Room temperature molten carboxylate salt hydrates. Synth. React. Inorg. Met.-Org. Chem. 2001, 31, 359−369. (12) Wang, G.; Hou, W.; Xiao, F.; Geng, J.; Wu, Y.; Zhang, Z. Lowviscosity triethylbutylammonium acetate as a task-specific ionic liquid for reversible CO2 absorption. J. Chem. Eng. Data 2011, 56, 1125− 1133. (13) Wang, G. N.; Dai, Y.; Hu, X. B.; Xiao, F.; Wu, Y. T.; Zhang, Z. B.; Zhou, Z. Novel ionic liquid analogs formed by triethylbutylammonium carboxylate-water mixtures for CO2 absorption. J. Mol. Liq. 2012, 168, 17−20. I


Effect of Temperature and Water Concentration on CO2 Absorption by

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