Effect of Alkyl Chain Length in Anions on Thermodynamic and

Article pubs.acs.org/IECR

Effect of Alkyl Chain Length in Anions on Thermodynamic and Surface Properties of 1-Butyl-3-methylimidazolium Carboxylate Ionic Liquids Airong Xu,†,‡ Jianji Wang,*,†,‡ Yajuan Zhang,† and Qingtai Chen† †

School of Chemistry and Environmental Science, Key Laboratory of Green Chemical Media and Reactions, Ministry of Education, Henan Normal University, Xinxiang, Henan 453007, P. R. China ‡ School of Chemical Engineering and Pharmaceutics, Henan University of Science and Technology, Luoyang, Henan 471003, P. R. China S Supporting Information *

ABSTRACT: Carboxylate-anion-based imidazolium ionic liquids (ILs) are powerful solvents for cellulose and lignin. However, little is known about their fundamental physicochemical properties. In this work, 1-butyl-3-methylimidazolium carboxylate ILs 1-butyl-3-methylimidazolium formate ([C4mim][HCOO]), acetate ([C4mim][CH3COO]), propionate ([C4mim][CH3CH2COO]), and butyrate ([C4mim][CH3(CH2)2COO]), in which the alkyl chain length in the anions is being varied in contrast to the more usual studies where alkyl chain length in the cations is varied, have been synthesized and their densities and surface tensions have been determined experimentally at different temperatures. By using these data, the molar volume, isobaric expansivity, standard entropy, lattice energy, surface excess entropy, vaporization enthalpy, and Hildebrand solubility parameter have been estimated for these ILs. From the analysis of structure−property relationship, the effect of alkyl chain length in the anions on these physicochemical properties of the ILs has been assessed and the dissolution of cellulose and lignin in these ILs has been discussed. Such knowledge is expected to be useful for understanding the nature of this class of solvent for the dissolution of biomacromolecules.

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INTRODUCTION Cellulose is the most abundant biorenewable and biodegradable resource in the world. Cellulose and its derivatives have been widely used in the areas such as fibers, tissues, papers, membranes, polymers, paints, and medical materials.1,2 However, due to the close packing by numerous inter- and intra- molecular hydrogen bonds, cellulose is extremely difficult to dissolve in water and in conventional organic solvents. This restricts the extensive application of cellulose, and the dissolution process of cellulose in the current use often suffers from environmental, energy, safety, or other problems.3−5 In recent years, it has been found that ionic liquids (ILs), which are a class of greener liquid materials which melt at temperatures below the boiling point of water, 1-butyl3-methylimidazolium chloride [C4mim]Cl,6 1-allyl-3-methylimidazolium chloride [Amim]Cl,7 1-ethyl-3-methylimidazolium formate [C2mim][HCOO],8 1-ethyl-3-methylimidazolium acetate [C2mim][CH3COO], and 1-butyl-3-methylimidazolium acetate [C4mim][CH3COO],9,10 can efficiently dissolve cellulose. Among these ILs, [C2mim][HCOO], [C2mim][CH3COO] and [C4mim][CH3COO] have attracted much attention due to their lower melting point, lower viscosity, lower corrosive character, and higher solubility of cellulose. However, it is surprising to find that, in spite of the importance of these ILs in dissolution of cellulose, little is still known about the fundamental physicochemical properties of this class of compounds, which are paramount for the design of any technological processes and are also important for better understanding of the nature of ILs. In this work, we have therefore synthesized the 1-butyl-3-methylimidazolium carboxylate ILs [C4mim][HCOO], [C4mim][CH3COO], [C4mim][CH3CH2COO], and [C4mim][CH3(CH2)2COO] in © 2012 American Chemical Society

which alkyl chain length in carboxylate anions varies (as shown in Figure 1). Densities and surface tensions of these ILs have been

Figure 1. Schematic structures for the investigated ILs.

determined experimentally as a function of temperature. Then, molar volume, isobaric expansivity, standard molar entropy, lattice energy, surface excess entropy, enthalpy of vaporization, and Hildebrand solubility parameter have been estimated for the ILs under study. These data were used to study the effect of alkyl chain length in the carboxylate anions on the physicochemical properties of the ILs, and to understand the nature of these ILs. We hope that the information provided here would be useful for the establishment of the structure−property relationship of ILs and for the molecular design of the ILs which are good solvents for the dissolution of biomacromolecules such as cellulose and lignin.

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EXPERIMENTAL SECTION Materials. 1-Methylimidazole (99%) was purchased from Shanghai Chem. Co. Lignin with a stated molecular weight of

Received: Revised: Accepted: Published: 3458

June 22, 2011 January 26, 2012 January 27, 2012 January 27, 2012 dx.doi.org/10.1021/ie201345t | Ind. Eng. Chem. Res. 2012, 51, 3458−3465

Industrial & Engineering Chemistry Research

Article

(m, 2H, CH2), 1.81 (t, 2H, CH2), 3.87 (s, 3H, NCH3), 4.19 (t, 2H, NCH2), 7.88 (s, 1H, NCH), 7.97 (s, 1H, NCH), and 10.31 (s, 1H, NCHN). Density Measurements. Densities of the ILs were measured using a vibrating-tube digital densimeter (DMA 60/602, Anton Paar, Austria) at the temperature range 298.15− 343.15 K. The temperature around the densimeter cell was controlled by circulating water from a constant-temperature bath (Schott, Germany). A CT-1450 temperature controller and a CK-100 ultracryostat were employed to maintain the bath temperature, and the temperature uncertainty is ±0.01 K. The densimeter was calibrated with known densities of pure water and dry air.12 The uncertainty in density was estimated to be ±1 × 10−5 g cm−3. Prior to each measurement, each dried IL was placed in a syringe pipette and then degassed in vacuum oven at 70 °C for 12 h in the presence of P2O5 to avoid possible a bubble which would generally result in unstable densimeter readings. In order to prevent absorption of water from the atmosphere, the ILs were then cooled to the experimental temperature in the vacuum oven under a dry nitrogen atmosphere. Surface Tension Measurements. Surface tensions of the ILs were measured with a DCA 315 tensiometer (Cahn Instruments) using a platinum plate (20 × 15 × 0.127 mm3) at the temperature range 298.15−343.15 K. The platinum plate was rinsed with water and ethanol and finally flame cleaned in a alcohol burner to eliminate any contaminants before every measurement. A 13-mL dried sample in a 50-mL beaker was found to be sufficient to wet the platinum plate and to prevent its interaction with the liquid meniscus, and the beaker was placed inside a thermostatic bath of the tensiometer sample chamber. The temperature around the beaker was controlled by circulating water from a HAAKE DC30-K20 temperature thermostat (Thermo Electron, Germany), and the temperature was maintained to be within ±0.1 K. All the measurements were performed under dry nitrogen atmosphere. Before measurement, each sample was allowed to attain thermal equilibrium for about 30 min at any given temperature setting. For each sample, at least three independent sets of immersion detachment cycles were measured, which allowed the determination of an average surface tension value. The tensiometer was calibrated with double distilled water according to the method provided by the manufacturer, and the uncertainty of the surface tension measurements was estimated to be about ±0.1 mJ m−2. Dissolution of Lignin in the ILs. Lignin sample was added into a 20 mL colorimetric tube containing 2.0 g of the dried IL. The tube was sealed with parafilm, and was then immersed in an oil bath (DF-101S, Gongyi Yingyu Instrument Factory). The uncertainty of the bath temperature was estimated to be ±0.5 °C. The mixture was stirred at 25 °C under argon atmosphere. Additional lignin was added until the solution became optically clear under polarization microscope (Nanjing Jiangnan Novel Optics Co. Ltd.). When lignin became saturated, judged by the fact that lignin could not be dissolved further within 2 h, its solubility (expressed by gram per 100 g of ionic liquid) could be calculated from the amount of the IL and lignin added.

28000 was purchased from Sigma Aldrich. 1-Bromobutane (98.0%) and anion exchange resin (Ambersep 900 OH) were from Alfa Aesar; acetic acid (>99.5%), formic acid (>88.0%), propanoic acid (>99.5%), and butyric acid (>99.5%) were from Shanghai Shiyi Chem. Reagent Co. Ltd. Ethyl acetate (>99.5%) was from Sinopharm Chem. Reagent Co. Ltd. Deuterated DMSO (DMSO-d6) used for NMR samples was purchased from Qingdao Weibo Tenglong Technol. Co. Ltd. These materials were used as received. Deionized water was doubly distilled over KMnO4, and the water with a conductivity of 1.2 × 10−6 S cm−1 was used throughout the experiments. Synthesis of the ILs. The studied ILs were synthesized in the following way. In order to prepare [C4mim]Br, the procedure described in the literature11 was closely followed. The reaction of 1-methylimidazole with an excess of 1-bromobutane was performed in 1,1,1-trichloroethane at 70 °C for 48 h. The solid product of [C4mim]Br was washed with 1,1,1-trichloroethane and then filtered. The residual solvent was removed by rotary evaporation, and the resulting product was dried under vacuum at 70 °C for 72 h in the presence of P2O5. 1H NMR (400 MHz, CDCl3, δ/ppm relative to TMS) data for [C4mim] Br are found to be the following: 0.85 (3H, t, but-CH3), 1.29 (2H, m, CH2), 1.81 (2H, m, CH2), 4.03 (3H, s, NCH3), 4.24 (2H, t, NCH2), 7.48 (1H, s, NCH), 7.61 (1H, s, NCH), and 10.23 (1H, s, NCHN). An aqueous solution of [C4mim]Br was allowed to pass through a column filled with anion exchange resin to obtain [C4mim][OH].8,10 In this process, the reaction was monitored by aqueous AgNO3. Once the substitution of bromide for hydroxide was completed, no precipitation of AgBr could be found. The aqueous [C4mim][OH] solution was then neutralized with an equal molar quantity of formic acid. After removing water by evaporation under reduced pressure, the viscous liquid [C4mim][HCOO] was thoroughly washed with diethyl ether, and finally dried under vacuum at 70 °C for 72 h in the presence of P2O5. The other ILs were prepared by a similar process as described for the preparation of [C4mim][HCOO]. The water content of the synthesized ILs was determined by Karl−Fisher titration using a 851 moisture titrator (Metrohm, Switzerland). It was found that the water content in [C4mim][HCOO], [C4mim][CH3COO], [C4mim][CH3CH2COO], and [C4mim][CH3(CH2)2COO] was 100.4, 99.1, 70.5, and 94.7 ppm, respectively. The 1H NMR data of the ILs are listed as follows. 1 H NMR (400 MHz, DMSO-d6, δ/ppm relative to TMS) for [C4mim][HCOO]: 0.84 (3H, t, but-CH3), 1.20 (2H, m, CH2), 1.73 (2H, m, CH2), 3.87 (3H, s, NCH3), 4.19 (2H, t, NCH2), 7.88 (1H, s, NCH), 7.90 (1H, s, NCH), 8.62 (1H, s, HCOO), and 9.99 (1H, s, NCHN). 1 H NMR (400 MHz, DMSO-d6, δ/ppm relative to TMS) for [C4mim][CH3COO]: 0.84 (3H, t, but-CH3), 1.18 (2H, m, CH2), 1.57 (3H, s, CH3CO2), 1.72 (2H, m, CH2), 3.87 (3H, s, NCH3), 4.19 (2H, t, NCH2), 7.90 (1H, s, NCH), 7.80 (1H, s, NCH), and 10.36 (1H, s, NCHN). 1 H NMR (400 MHz, DMSO-d6, δ/ppm relative to TMS) for [C4mim][CH3CH2COO]): 0.82 (t, 3H, but-CH3), 0.85 (t, 3H, ethyl-CH3), 1.18 (m, 2H, CH2), 1.72 (m, 2H, CH2), 1.82 (m, 2H, CH2), 3.88 (s, 3H, NCH3), 4.19 (t, 2H, NCH2), 7.90 (s, 1H, NCH), 7.99 (s, 1H, NCH), and 10.41 (s, 1H, NCHN). 1 H NMR (400 MHz, DMSO-d6, δ/ppm relative to TMS) for [C4mim][CH3(CH2)2COO]: : 0.76 (t, 3H, but-CH3), 0.83 (t, 3H, propyl-CH3), 1.18 (m, 2H, CH2), 1.39 (m, 2H, CH2), 1.72

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RESULTS AND DISCUSSION Molar Volume, Standard Entropy, Lattice Energy, and Isobaric Expansivity. Table 1 shows the experimental densities of [C4mim][HCOO], [C4mim][CH3COO], [C4mim][CH3CH2COO], and [C4mim][CH3(CH2)2COO] over the temperture range from 298.15 to 343.15 K at atmospheric pressure. 3459

dx.doi.org/10.1021/ie201345t | Ind. Eng. Chem. Res. 2012, 51, 3458−3465


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Table 1. Density, ρ, of the Investigated ILs at Different Temperatures ρ (g cm−3) IL

298.15 K

303.15 K

313.15 K

323.15 K

333.15 K

343.15 K

[C4mim][HCOO] [C4mim][CH3COO] [C4mim][CH3CH2COO] [C4mim][CH3(CH2)2COO]

1.07341 1.05223 1.03491 1.03222

1.07037 1.04659 1.03194 1.02890

1.06441 1.04049 1.02593 1.02274

1.05854 1.03445 1.01983 1.01637

1.05273 1.02839 1.01368 1.01002

1.04689 1.02242 1.00755 1.00371

where M is the molar mass of the ILs. The calculated Vm values of the ILs are listed in Table 3. It can be seen that, for the given cation [C4mim]+, the molar volumes of the ILs increase with the effective anion size in the order: [HCOO]− < [CH3COO]− < [CH3CH2COO]− < [CH3(CH2)2COO]−. Actually, a linear relationship was observed between molar volume of the ILs and number of C atoms in alkyl chain of the carboxylate anions at any given temperature. Figure 3 (the left-hand side) shows such a linear plot at 298.15 K with a slope of 16.0 cm3 mol−1, indicating that a regular addition per CH2 group in the anion alkyl chain of the ILs contributes to an increase of 16.0 cm3 mol−1 in their molar volumes. This value is close to the CH2 contribution of 16.6 cm3 mol−1 from 1-alkyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide ([Cnmim][NTf2]), 16.7 cm3 mol−1 from 1-alkyl-3-methylimidazolium alanine ([Cnmim][Ala]), 16.9 cm3 mol−1 from n-alcohols, 16.4 cm3 mol−1 from n-amines, and 16.1 cm3 mol−1 from n-paraffins.16,20,21 This suggests that CH2 contribution to the molar volume is independent of the nature of both anions and cations of ILs, and can be predicted from homologues of commonly used organic compounds. As suggested by Yang and co-workers,20 the relationship between standard entropy and molecular volume at 298 K established by Glasser16

Figure 2. Temperature dependent on density of the ILs: (a) [C4mim][HCOO]; (b) [C4mim][CH3COO]; (c) [C4mim][CH3CH2COO]; (d) [C4mim][CH3(CH2)2COO].

The temperature dependence of the ILs densities is shown in Figure 2. An essentially linear decrease in density was observed as the temperature increases for all the ILs investigated. At the same time, it is noted that densities of the ILs decrease with an increase of the alkyl chain length in the carboxylate anions (Table 1), which indicates that the anionic structure has an important impact on the densities of these ILs. A similar trend was previously reported for the ILs with different alkyl chain lengths in cations.13−16 Usually, liquid densities at 0.1 MPa can be correlated with the Tait equation:17−19 ρ(g cm−3) = a1 + a2T + a3T 2

S° (J K −1 mol−1) ≈ 1246.5V (nm3) + 29.5

can be used to calculate the standard entropy at 298.15 K from the molar volume data of the ILs, and the result is presented in Table 4. In eq 3, V is molecular volume of the ILs and can be calculated from V (nm3) = Vm (nm3 mol−1)/NA, where the Avogadro’s constant NA = 6.02245 × 1023 molecule mol−1. From Figure 3 (the right-hand side), it can be seen that the standard entropy of the ILs increases linearly with increasing number of C atoms in alkyl chain of the carboxylate anions at 298.15 K, suggesting the less organization in the ILs as alkyl chain length in carboxylate anions increases. Actually, this linear relationship can be predicted from eq 3 and the linear relationship between molar volume of the ILs and number of C atoms in alkyl chain of their anions shown in the left-hand side of Figure 3. It is found that the slope of this straight line is 33.0 J K−1 mol−1, which is in reasonable agreement with the value of 34.6 J K−1 mol−1 from [Cnmim][Ala],20 33.9 J K−1 mol−1 from [Cnmim][BF4], 35.1 J K−1 mol−1 from [Cnmim][NTf2],16 and 32.2 J K−1 mol−1 from an extended group of organic compounds.22 This clearly indicates that a regular addition of

(1)

Our density data have been fitted to this equation, and the coefficients a1, a2, and a3 obtained by a least-squares analysis are given in Table 2. It can be seen that the Tait equation correlates well with our IL density data, and the coefficient a3 is so small that a3T2 in eq 1 can be neglected. Therefore, the Tait equation becomes linear under these circumstances, which is consistent with the result shown in Figure 2. From the experimental density data, the molar volume of the ILs, Vm, was calculated using the equation Vm(cm3mol−1) = M /ρ(g mol−1/g cm−3)

(3)

(2)

Table 2. Coefficients of Equation 1 and Their Standard Deviation IL

a1 (g cm−3)

a2 × 104 (g cm−3 K−1)

a3 × 107 (g cm−3 K−2)

SD × 105 (g cm−3)


1.27042 1.24765 1.19541 1.19941

−7.2489 −7.1519 −4.7642 −5.0174

2.1429 1.7143 −2.0714 −2.0000

2.7 1.7 2.5 5.9

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Table 3. Molar Mass, M, and Molar Volume, Vm, of the Investigated ILs at Different Temperatures Vm (cm3 mol−1) −1

IL

M (g mol )

298.15 K

303.15 K

313.15 K

323.15 K

333.15 K

343.15 K


184.24 198.26 212.29 226.32

171.64 188.42 205.13 219.26

172.13 189.43 205.72 219.96

173.09 190.55 206.92 221.29

174.05 191.66 208.16 222.67

175.01 192.79 209.42 224.07

175.99 193.91 210.70 225.48

have been used to calculate thermodynamic properties of some ionic liquids.26,27 The crystal energies for the ionic liquids calculated from eq 4 are included in Table 4. It is found that the values of crystal energy for the carboxylate ILs are smaller than those of inorganic fused salts. For example, the crystal energy is 613 kJ mol−1 for the fused CsI,28 which is about 12% higher than that of [C4mim][CH3(CH2)2COO]. This low crystal energy is the underlying reason for the formation of liquid salts at room temperature. In addition, the values of crystal energy decrease with increasing number of C atoms in alkyl chain of the carboxylate anions. A similar trend was reported for [Cnmim][Ala] which have different alkyl chain lengths in cations.20 By using the density data at different temperatures, the isobaric expansivity, αp, of the ILs has been calculated from the following equation:

Figure 3. Linear relationship between the molar volume or the standard entropy of the ILs and the C atom number (nc) in alkyl chain of the carboxylate anions at 298.15 K.

α p (K −1) =

Table 4. Standard Entropy and Lattice Energy for the Investigated ILs at 298.15 K

(5)

IL

V (nm3)

S° (J K−1 mol−1)

UPOT (kJ mol−1)


0.2851 0.3130 0.3408 0.3642

384.9 419.7 454.3 483.5

567.7 560.0 553.1 547.9

The data in Table 5 indicate that αp values of the ILs are smaller than those of most organic liquids and similar to that of water, confirming the ionic nature of the ILs. At the same time, we found that the isobaric expansivity increases with increasing alkyl chain length of the carboxylate anions at a given temperature. This implies that degree of ordering in the ILs decreases with increasing alkyl chain length of the anions, which supports our result obtained from the analysis of standard entropies of the ILs. Surface Tension and Surface Excess Entropy. For the ILs studied, the measured surface tension data at different temperatures are reported in Table 6, and the temperature dependence of the surface tensions is shown in Figure 4. It can be seen that the surface tension of a given IL decreases linearly with a rise in temperature, and the increase of the alkyl chain length in the carboxylate anions results in a lowering of the surface tension of the ILs at a given temperature. The latter trend may be explained by Langmuir’s principle of independent surface action.29 On the basis of this principle, surface tension is related to the part of the molecule that is actually present at the interface. Since anions and cations are present at the surface of ILs,21,30 both of them should contribute to the surface free energy. Therefore, the application of this principle needs to know the surface tensions of the separated ions. Although these

each CH2 group in the alkyl chain of the anions contributes to an increase of 33.0 J K−1 mol−1 in the standard entropy of the carboxylate ILs, and the value can also be approximately predicted from the data of [Cnmim]+ based ILs and some organic compounds. For simple salt of type MX, the crystal energy UPOT is given by23,24 UPOT (kJ mol−1) ≈ 2I {α(V (nm3))−1/3 + β}

⎛ ∂ln ρ ⎞ 1 ⎛ ∂Vm ⎞ 1 ⎛ ∂ρ ⎞ ⎟ ⎜ ⎟ = − ⎜ ⎟ = −⎜ ⎝ ∂T ⎠P Vm ⎝ ∂T ⎠P ρ ⎝ ∂T ⎠P

(4)

where α and β are fitted coefficients of the salts, I is the ionic strength, and V is the molecular volume (in nm3). For 1:1 inorganic salts, I = 1, α = 117.2 kJ mol−1nm, and β = 51.9 kJ mol−1. However, Gutowski et al.25 noted that, for ionic liquids which are 1:1 salts but contain organic type cations, the α and β coefficients need to be revised as α = 83.3 kJ mol−1 nm, and β = 157.3 kJ mol−1. Very recently, these revised coefficients

Table 5. Isobaric Expansivity, αp, Data of the Investigated ILs at Different Temperatures αp × 104 (K−1) IL

303.15 K

313.15 K

323.15 K

333.15 K

343.15 K


5.5621 5.8405 5.8356 6.0482

5.5515 5.8418 5.9095 6.1252

5.5405 5.8429 5.9845 6.2034

5.5292 5.8438 6.0607 6.2829

5.5175 5.8445 6.1380 6.3636

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Table 6. Values of the Surface Tension, γ, for the Investigated ILs at Different Temperatures γ (mJ m−2) IL

298.15 K

303.15 K

313.15 K

323.15 K

333.15 K

343.15 K


45.9 39.2 36.2 34.3

45.5 39.0 35.8 33.9

44.7 38.7 35.4 33.2

44.0 38.3 35.1 32.8

43.4 38.1 34.8 32.1

42.8 37.7 34.4 31.6

[C4mim][CH3COO], [C4mim][CH3CH2COO], and [C4mim][CH3(CH2)2COO] also reduce in the same sequence. The surface excess energy Es and surface excess entropy Ss of the ILs can be derived from the measured surface tension data according to the following equation:29 γ (mJ m−2) = Es − TSs

Obviously, Es (mJ m−2) and Ss (mJ m−2 K−1) values can be obtained, respectively, from intercept and slope of the straight lines shown in Figure 4. Thus calculated values are presented in Table 8. It can be seen that surface excess entropies of the ILs are remarkably lower than that of water, benzene, and pyridine33 (0.138, 0.13, and 0.1369) mJ m−2 K−1, respectively. This indicates that a high degree of organization is present in the ILs, which is in agreement with the results reported by Santos et al.34 Additionally, it is also noted that the surface excess entropies are alkyl chain length dependent, and they increase as the alkyl chain in the carboxylate anions is lengthened. For example, with the addition of methylene group to the carboxylate anion, the surface excess entropy increases in the following series: [C4mim][CH3COO] (0.032 mJ m−2 K−1) < [C4mim][CH3CH2COO] (0.036 mJ m−2 K−1) < [C4mim][CH3(CH2)2COO] (0.057 mJ m−2 K−1) (Table 8). This indicates that organization of the ILs reduces as C atom number in the alkyl chain of the carboxylate anions increases from 1 to 3. However, it should be pointed out that the surface excess entropy of [C4mim][HCOO] (0.067 mJ m−2 K−1) is exceptionally higher than that of [C4mim][CH3COO] and [C4mim][CH3CH2COO]. We cannot explain this result at the present stage. On the other hand, the data in Table 8 clearly indicate that the surface excess energies of the investigated ILs are about one-third of the fused salts like NaNO3 (146 mJ m−2). This suggests that the interaction energy between cations and anions in the investigated ILs is much lower than that in inorganic fused salts.20 Enthalpy of Vaporization and Hildebrand Solubility Parameter. One of the unique properties for room temperature ILs is negligibly low vapor pressure, which means that direct experimental determination of the molar enthalpy of vaporization, ΔlgHm°, is very difficult for these compounds. However, it is possible for us to estimate Δ1gHm° indirectly

Figure 4. Temperature dependent on surface tension of the ILs: (a) [C 4 mim][HCOO]; (b) [C 4 mim][CH 3 COO]; (c) [C 4 mim][CH3CH2COO]; (d) [C4mim][CH3(CH2)2COO].

Table 7. Values of the Surface Tension, γ, of Some Compounds at 298.15 K26,27 entry

compd

γ (mJ m−2)

1 2 3 4 5

imidazole formic acid acetic acid propionic acid butyric acid

43 38.17 27.08 26.15 26.19

(6)

are not feasible, extrapolated values can be employed to make a qualitative comparison. Table 7 summarizes the surface tensions of some compounds whose molecules are similar to the parts of the IL cations and anions.31,32 Similarity of surface tension of the ILs studied with the average value of the surface tension of imidazole and carboxylic acid is a clear indication of the validity of Langmuir’s principle. It is expected that if the surface tension data of butylimidazole or methylimidazole are used, more reasonable agreement can be seen. Unfortunately, no such data have been reported in the literature. In addition, considering the fact that the surface tension of carboxylic acids decreases in the series HCOOH > CH3COOH > CH3CH2COOH > CH3(CH2)2COOH (Table 7), it is easy to understand why the surface tensions of [C4mim][HCOO],

Table 8. Surface Excess Entropy Ss, Surface Excess Energy Es, Enthalpy of Vaporization Δ1gHm° and Hildebrand Solubility Parameter δH of the Investigated ILs at 298.15 K IL [C4mim][HCOO] [C4mim][CH3COO] [C4mim][CH3CH2COO] [C4mim][CH3(CH2)2COO]

Es (mJ m−2) 65.7 48.7 46.7 51.1

± ± ± ±

0.8 0.5 0.6 0.6

Ss (mJ m−2 K−1)

ΔlgHm° (kJ mol−1)

δH ((MPa)0.5)

± ± ± ±

95.2 124.0 121.6 120.5

23.2 25.4 24.1 23.2

0.067 0.032 0.036 0.057 3462

0.003 0.002 0.002 0.003



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of lignin,40 and thus [C4mim][CH3CH2COO] is expected to be a good solvent for the dissolution of lignin. In order to verify this prediction, we determined the solubility of lignin in the ionic liquids experimentally at 298 K. Our results showed that the solubility of lignin in the ionic liquids decreases in the following order: [C4mim][CH3CH2COO] (11 g per 100 g of IL) > [C4mim][CH3COO] (6 g per 100 g of IL) > [C4mim][HCOO] (5 g per 100 g of IL) = [C4mim][CH3(CH2)2COO] (5 g per 100 g of IL). In addition, it was reported that [C4mim][CH3COO] is a better solvent for cellulose than [C4mim][HCOO].11 Although no δH value has been reported for cellulose so far in the literature, it can be expected that the δH value of cellulose would be close to or greater than 25.4 (MPa)0.5 which is the δH value of [C4mim][CH3COO] estimated in the present work.

from the empirical equation put forward by Kabo and Verevkin35 g

Δ1 Hm° (kJ mol−1) = A(γVm 2/3NA1/3) + B −2

(7)

−1

where γ (J m ), Vm (m mol ), and NA are, respectively, the surface tension in J m−2, molar volume in m3 mol−1, and Avogadro’s constant, and A (=0.01121) and B (=2.4 kJ mol−1) at 298 K are empirical parameters. The values of molar enthalpy of vaporization for the ILs calculated from eq 7 are given in Table 8. Given that solubility parameters are useful for the prediction of solubility of a solute in a solvent,36 the Hildebrand’s solubility parameters (δH's) for the investigated ILs have been calculated from the equation36 3

δ H ((MPa)0.5 ) = [(Δ1g Hm° − RT )/Vm]0.5

(8)

where ΔlgHm° (J mol−1) and Vm (cm3 mol−1) are, respectively, the vaporization enthalpy in J mol−1 and molar volume in cm3 mol−1 for the ILs, and R and T have their usual meaning. The results at 298.15 K calculated by using the values of molar volume and vaporization enthalpy are also included in Table 8. To confirm the reliability of the δH values obtained here, we estimated the ΔlgHm° and Vm values of [C8mim][BF4] from its surface tension37 and density38 data at 298.15 K, respectively. These values were then used to calculate the δH value of [C8mim][BF4] by eq 8. The obtained δH value (22.0 (MPa)0.5) for this IL is very close to the experimental value (22.5 (MPa)0.5) reported in the literature,39 indicating that the δH value provided in this study is reasonable. It can be seen from Table 8 that the δH values increase in the following order: [C4mim][HCOO] = [C4mim][CH3(CH2)2COO] < [C4mim][CH3CH2COO] < [C4mim][CH3COO]. This is similar to the order of the enthalpy of vaporization for these ILs. It was reported that 1H NMR chemical shift (δ) magnitude of the proton in the 2-position of the imidazolium ring can be used as a measure for the capacity of anion of ILs to form hydrogen bonds with their cations.10 Therefore, we determined the 1H NMR chemical shifts of the ILs relative to TMS in DMSO-d6 at the concentration of 1.0 mol kg−1 in order to see what kind of interaction contributes to the order observed above. It is found that the capacity of the anions to form hydrogen bonds with their cations follows the order: [HCOO]− (δ = 9.989 ppm) < [CH3(CH2)2COO]− (δ = 10.311 ppm) < [CH3COO]− (δ = 10.361 ppm) < [CH3CH2COO]− (δ = 10.405 ppm). This order suggests that the presence of alkyl chains which are electron-donating groups promotes the formation of hydrogen bonds between O− of the anions and proton in the 2-position of the imidazolium cations. However, as the number of C atoms in the alkyl chains is equal to 3, the steric hindrance effect of the alkyl chains began to reduce the electron-donating effect for the formation of hydrogen bonds. A comparison of the order in hydrogen bond interaction with that in solubility parameter observed reveals that the hydrogen bond interaction does contribute greatly to the order observed, but it is not the only contribution. It is well-known that δH value is effectively a measure of the strength of molecular interactions between solvent molecules. These parameters can be used to predict the solubility of biomacromolecules in the ILs, and the maximum solubility is observed when the δH values of the biomacromolecule and the IL are identical.39 For example, the δH value (24.1 (MPa)0.5) of [C4mim][CH3CH2COO] is very close to that (24.6 (MPa)0.5)

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CONCLUSIONS

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ASSOCIATED CONTENT

In the present work, values of density, surface tension, molar volume, isobaric expansivity, standard entropy, lattice energy, surface excess entropy, enthalpy of vaporization, and Hildebrand solubility parameter of [C4mim][HCOO], [C4mim][CH 3 COO], [C 4 mim][CH 3 CH 2 COO], and [C 4 mim][CH3(CH2)2COO] have been estimated in order to assist the possible industrial application of these ILs in dissolution of cellulose and lignin. It is found that, at a given temperature, the anionic structure has an important impact on the physicochemical properties of these ILs. The molar volumes and standard entropies increase linearly with the number of CH2 in the alkyl chain of the anions. Also, the values of surface tensions, isobaric expansivities, and surface excess entropies increase, but crystal energy decreases with the increase in the number of C atoms in alkyl chain of the carboxylate anions. However, among the ILs studied, [C4mim][CH3COO] has the maximum enthalpy of vaporization and Hildebrand solubility parameter. It is also interesting to find that, at a given temperature, the contribution per CH2 group in the anions to the molar volumes and the standard entropy is independent of the nature of both anions and cations of the ILs, and can be predicted from homologues of commonly used organic compounds. The lower crystal energy of these ILs is the essential reason for their formation of liquid salts at room temperature, and due to the ionic nature of these ILs, their surface excess entropies are only about one-third of that for most organic liquids and water. In addition, the higher solubility of cellulose in [C4mim][CH3COO] compared to [C4mim][HCOO] has been ascribed to the higher solubility parameter of [C4mim][CH3COO], and [C4mim][CH3CH2COO] is expected to be a good solvent for the dissolution of lignin. The information provided here would be useful for the establishment of the structure−property relationship of ILs, and may provide helpful clues for better understanding the nature of this class of solvent for the dissolution of biomacromolecules and for the molecular design of the ILs with desirable properties.

S Supporting Information *

Additional figure. This material is available free of charge via the Internet at http://pubs.acs.org. 3463



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Article

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported financially by the National Basic Research Program of China (973 Program, No. 2011CB211702), and the Innovation Scientists and Technicians Troop Construction Projects of Henan Province (No. 092101510300).

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Effect of Alkyl Chain Length in Anions on Thermodynamic and

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