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Solubilities of p-coumaric and caffeic acid in ionic liquids and organic solvents
J. Chem. Thermodynamics 62 (2013) 69–78
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Solubilities of p-coumaric and caffeic acid in ionic liquids and organic solvents Efthimia I. Alevizou, Epaminondas C. Voutsas ⇑ Laboratory of Thermodynamics and Transport Phenomena, School of Chemical Engineering, National Technical University of Athens, 9 Heroon Polytechniou Str., Zographou Campus, 15780 Athens, Greece
a r t i c l e
i n f o
Article history: Received 24 July 2012 Received in revised form 21 February 2013 Accepted 23 February 2013 Available online 4 March 2013 Keywords: Antioxidants p-Coumaric acid Caffeic acid Ionic liquids Solubility measurements Correlation
a b s t r a c t The solubilities of two cinnamic acid derivatives, namely p-coumaric acid and caffeic acid, in six 1alkyl-3-methyl imidazolium based ionic liquids composed of the PF6, BF4, TFO and TF2N anions, and in two organic solvents, t-pentanol and ethyl acetate, have been measured at the temperature range of about (303 to 317) K. The p-coumaric acid was found to be more soluble than caffeic acid in all studied solvents. Higher solubilities of both acids were observed in the ionic liquids composed of the BF4 and TFO anions. The increase of the alkyl chain length on the cation invokes a decrease in solubility in the case of hydrophilic ionic liquids composed of BF4 anion, while in the case of hydrophobic ones composed of PF6 anion an increase in the solubility is observed. Between the two organic solvents t-pentanol is better solvent than ethyl acetate for both acids. Moreover, using the van’t Hoff equations the apparent Gibbs energy, enthalpy, and entropy of solution were calculated. Finally, successful correlation of the experimental data was achieved with the UNIQUAC and the NRTL activity coefficient models, while poor predictions of the solubility of the two acids in the organic solvents were obtained with two UNIFAC models. Ó 2013 Elsevier Ltd. All rights reserved.
1. Introduction Cinnamic acid derivatives (CADs), such as p-coumaric acid and caffeic acid, are natural antioxidants belonging to the family of phenolic acids and are widely distributed throughout the plant kingdom. Fruits, vegetables, spices, aromatic herbs, cereals, coffee beans and beverages, olives, propolis, sunflower, barks and wine are natural sources of these compounds [1–4]. Due to their almost universal distribution, they are an integral part of the human diet. Epidemiological studies suggest a link between the consumption of whole grain products containing CADs and prevention of chronic, degenerative diseases associated with oxidative damage (namely coronary heart diseases, cardiovascular diseases, diabetes, arthritis, cataract formation, aging and cancer) [5,6]. Many of the health protective effects of phenolic compounds have been ascribed to their antioxidant, antimutagenic, anti-inflammatory, antimicrobial, antiviral, antialergic, immunoprotective, ultraviolet (UV) filtering properties and other biological and pharmacological activities Abbreviations: CADs, cinnamic acid derivatives; ILs, ionic liquids; p-CA, p-coumaric acid; CA, caffeic acid; bmim, 1-butyl-3-methyl imidazolium; omim, 1-octyl-3-methyl imidazolium; PF6, hexafluorophosphate; BF4, tetrafluoroborate; TF2N, bis(trifluoromethanesulfonyl)imide; TFO, trifluromethanesulfonate; HPLC, high-performance liquid chromatography; DSC, differential scanning calorimetry; TgA, thermogravimetric analysis; calc, calculated; exp, experimental; SD, standard deviation; AARE, absolute average relative error; ARE, absolute relative error. ⇑ Corresponding author. Tel.: +30 210 772 3971; fax: +30 210 772 3155. E-mail address: [email protected] (E.C. Voutsas). 0021-9614/$ - see front matter Ó 2013 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.jct.2013.02.013
[7,8]. The above benefits in combination with a growing concern about the safety of synthetic antioxidants [9] make CADs high value raw materials for the synthesis of different molecules with industrial interest, such as drugs, cosmetics, antiseptics and flavors [10]. Common problems encountered in chemical synthesis, recovery and separation processes, such as the use of toxic, volatile and flammable organic solvents as well as difficulties in product purification can be handled with the use of ionic liquids (ILs); a new class of solvents that can lead to more environmentally friendly applications and processes as compared to those where conventional organic solvents are used. This characteristic arises mainly from their negligible vapor pressures at room temperature and therefore low volatility and flammability, which minimizes health and safety risk in industry and the chance of loss to atmosphere making their recycling and reusability feasible. Of course, the development of cost and energy efficient technologies for solute recovery from ionic liquids is a challenge towards their industrial application. Ionic liquids are molten salts, consisting of large asymmetric organic cations and organic or inorganic anions. The low symmetry, high vibrational freedom and charge delocalization of the ions composing an IL reduce the stability of the crystalline phases, and thus their melting temperatures [11]. Unlike molecular liquids, their ionic nature results in a unique combination of distinctive properties, such as high thermal stabilities, large liquidus range and high solvating capacity for organic, inorganic and
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organometallic compounds [12,13]. The last is one of the most attractive features of ILs especially for recovery and separation processes, as well as final product purification. It results from ILs’ ability to be tailormade for a specific purpose by careful selection among a huge diversity of cations and anions, or substitutes to the cation. Therefore, ILs are often referred to as ‘‘designer solvents’’ [14]. For the design and optimization of processes where CADs are involved, suitable solvents – either classical organic ones or ionic liquids – have to be selected. The successful completion of such a task requires the knowledge of reliable phase equilibrium data. With reference to solubilities of antioxidants in organic solvents, there are some data in the literature, such as solubility of flavonoids in acetonitrile, acetone and tert-amyl alcohol [15], or solubility of luteolin in methanol, ethanol, 1-propanol, 2-propanol, 1-butanol, acetone, hexane, and DMSO [16], but for phenolic acids only their solubility in water has been measured [8,17]. Concerning ILs, although much experimental work has been conducted so far on measurements of thermophysical properties of the pure ILs and phase equilibrium in systems containing ILs, solubility data of antioxidants in ILs are very scarce [18–20]. Some other works have also reported solubility data in ILs of antibiotics [21], dibasic carboxylic acids [22], drugs [23–25] and inorganic salts [26]. In this work, the solubilities of two CADs, namely p-coumaric acid and caffeic acid, were measured in six imidazolium based ILs composed of the PF6, BF4, TF2N and TFO anions and they were compared to those measured in two organic solvents, t-pentanol and ethyl acetate. In addition, the normal melting point temperature and the enthalpy of fusion of the two acids were determined by differential scanning calorimentry and their possible decomposition upon melting was tested by thermogravimetric analysis. Moreover, by utilizing the solubility data and the van’t Hoff equations, the apparent thermodynamic functions relative to the solution of the two CADs have been determined. Finally, the capability of the UNIQUAC and the NRTL activity coefficient models to correlate the solubility data was tested and two UNIFAC versions were used to predict the solubilities of the two acids in the organic solvents.
2. Experimental section 2.1. Materials p-Coumaric acid (p-CA; purity > 98%; CAS No. 501-98-4; C9H8O3) was purchased from Sigma Aldrich Co. Caffeic acid (CA, purity > 98% CAS No. 331-39-5; C9H8O4) was purchased from Acros
Organics. 2-methyl-2-butanol (t-pentanol; purity > 99%; CAS No. 75-85-4; C5H12O) was purchased from Sigma Aldrich Co. Ethyl acetate (purity > 99.8%; CAS No. 141-78-6; C4H8O2) was purchased from Merck KGaA. 1-butyl-3-methyl-imidazolium hexafluorophosphate (bmimPF6; purity > 98%; CAS No. 174501-64-5; C8H15F6N2P), 1-octyl-3-methyl-imidazolium hexafluorophosphate (omimPF6; purity > 98%; CAS No. 304680-36-2; C12H23F6PN2), 1-butyl-3methyl-imidazolium tetrafluoroborate (bmimBF4; purity > 98%, CAS No. 174501-65-6; C8H15BF4N2), 1-octyl-3-methyl-imidazolium tetrafluoroborate (omimBF4; purity > 98%; CAS No. 244193-52-0; C12H23BF4N2), 1-butyl-3-methyl-imidazolium bis(trifluoromethanesulfonyl)imide (bmimTF2N; purity > 98%, CAS No. 17489983-3; C10H15F6N3O4S2), 1-butyl-3-methyl-imidazolium trifluoromethanesulfonate (bmimTFO, purity > 98%, CAS No. 174899-66-2; C9H15F3N2O3S) were purchased from Solchemar Lda., Portugal. Methanol (99.9% purity) and acetonitrile (99.9% purity) were purchased from Fisher Scientific. Water (HPLC gradient) was purchased from Merck. Table 1 summarizes the source and purities of compounds used in this work. The molecular structures of the investigated cinnamic acid derivatives (CADs) and the cations and anions of the ionic liquids (ILs) are presented in figure 1. 2.2. Solubility measurements and procedure The equilibrium runs were performed in a Thermomixer Comfort (Eppendorf). This apparatus combines stirring up to 1400 rpm and heating (up to 99 °C) or cooling (13 °C below room temperature). The temperature readings of the thermomixer have been checked using an external thermometer calibrated with the Isocal-6 VenusPlus 2140 calibrator of temperature sensors. The accompanied thermoblock can accommodate 24 micro test tubes of 2 mL each. Approximately 200 lL of each solvent were placed in the 2 mL volume safe-lock micro test tube. An excess quantity of each acid was then added. The tubes were placed in the apparatus, which was set to the desired temperature measured with an uncertainty of ±0.1 K, and were stirred at 1400 rpm until equilibrium was reached. Based on Mota et al.’s solubility measurements of caffeic acid in water [8], where a quite long shaking time was needed (64 to 84) h, it was concluded that in our case of indirect stirring and more viscous solvents the equilibrium times would be even longer. To determine the time required to reach equilibrium, a preliminary measurement of the solubility of the most soluble acid (p-CA) in the most viscous IL omimBF4 was performed. Samples were taken and analyzed after (24, 72, 120, 144, 168, and 192) h of stirring. It was found that the difference among the solubility values measured after (144, 168, and 192) h was less than 2%,
TABLE 1 List of compounds, their source and purities.
a b
Chemical name
Source
Purity %
3-(4-Hydroxyphenyl)-2-propenoic acid a 3-(3,4-Dihydroxyphenyl)-2-propenoic acid b 1-Butyl-3-methyl-imidazolium hexafluorophosphate 1-Octyl-3-methyl-imidazolium hexafluorophosphate 1-Butyl-3-methyl-imidazolium tetrafluoroborate 1-Octyl-3-methyl-imidazolium tetrafluoroborate 1-Butyl-3-methyl-imidazolium bis(trifluoromethanesulfonyl)imide 1-Butyl-3-methyl-imidazolium trifluoromethanesulfonate 2-Methyl-2-butanol Ethyl acetate Methanol Acetonitrile Water
Sigma Aldrich Co Acros Organics Solchemar Lda Solchemar Lda Solchemar Lda Solchemar Lda Solchemar Lda Solchemar Lda Sigma Aldrich Co Merck KGaA Fisher Scientific Fisher Scientific Merck KGaA
98.0 98.0 98.0 98.0 98.0 98.0 98.0 98.0 99.0 99.8 99.9 99.9 99.99
p-Coumaric acid. Caffeic acid.
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OH
OH OH
OH
OH
O
p-coumaric acid (p-CA)
O
caffeic acid (CA) +
+
N
N
N
N
1-butyl-3-methyl imidazolium cation (bmim+) 1-octyl-3-methyl imidazolium cation (omim+)
hexafluorophosphate anion (PF6-)
bis(trifluoromethylsulfonyl)imide anion (TF2N-)
tetrafluoroborate anion (BF4-)
trifluoromethylsulfonate anion (TFO-)
FIGURE 1. Structures of investigated cinnamic acid derivatives and cations/anions of ionic liquids.
which led to the conclusion that a time frame of 144 h is sufficient for achieving equilibrium for all solutions. After equilibrium was reached, solution samples were withdrawn using adjustable micropipettes. The sampling tips of the pipettes were preheated at the temperature of the experiment to avoid solid precipitation. Before sampling, each solution was centrifuged for a short time (60 to 90) s to enhance the physical separation of the two phases, liquid and solid, and to ensure that no precipitated solute was also drawn away during sampling because of the high viscosities of the ILs. All experiments were performed in triplicate. The water content in the studied solvents was measured by Karl-Fisher titration (volumetric titrator, TitroLine KF, SCHOTT Instruments) and was found to be: 0.97% w/w for bmimBF4, 1.45% for omimBF4, 1.35% for bmimPF6, 1.4% for omimPF6, 0.98% for bmimTF2N, 1.55% for bmimTFO, 0.06% for t-pentanol and 0.14% for ethyl acetate. The effect of the water traces in the ILs on the solubility was also investigated. For this purpose, solubility measurements were performed using dried (at 80 °C under vacuum for 48 h) ILs, and when they were compared with those obtained in the ILs used as purchased, no difference was found. For solubility measurements all solvents were used as received. Samples were analyzed by reverse phase high-performance liquid chromatography (HPLC). The HPLC system was equipped with a Jusco PU-1580 pump, an ICI LC 1200 UV/vis detector set at 320 nm, and a BDS Hypersil C18 column ((250 4.6) mm2, particle size 5 lm, Thermo Scientific). The eluent solution used was a mixture of acetonitrile/water (80/20 v/v) with 0.1% acetic acid. The flow rate was adjusted to 1 mL min1, the column temperature was kept constant at 30 °C, and the injection volume was 50 lL. The eluent solution was degassed with Helium.
2.3. Differential scanning calorimetry (DSC) and thermogravimetric analysis (TgA) For the calculation of the solid solubility, the normal melting point temperature (Tm) and heat of fusion (DfusH) of p-coumaric and caffeic acid are needed. Although there are such data in the literature, they show notable differences mostly regarding the melting point values and whether or not the two acids decompose upon melting. For example, Park et al. [27] have determined the melting point temperature of caffeic acid (Tm = 232.32 °C) with a DSC analysis without referring to a possible decomposition, while Mota et al. [8] could not determine it – also with a DSC analysis – because of its decomposition upon melting. The NIST Chemistry Webbook [28] reports a melting temperature range of (202 to 216) °C for p-coumaric acid. Saldana et al. [29] present the Tm of p-coumaric acid to be 214 °C and that of caffeic acid to be 196 °C, both taken from the Material Safety Data Sheets of Alfa and Sigma–Aldrich, respectively. The same values are presented also by Murga et al. [30] taken from Sigma–Aldrich. It is therefore concluded that there is a high uncertainty for the normal melting point temperatures of the two CADs. Thus, it was necessary to carry out a DSC analysis in combination with a TgA one, so as to determine the Tm and DfusH of the studied materials and also to investigate possible decomposition upon melting. For DSC analysis a Mettler – Toledo DSC1 apparatus was used, the temperature range was (50 to 300) °C with a heating rate of 10 °C/min under nitrogen flow. DSC apparatus was calibrated with indium, zinc and n-heptane at the onset temperature. For TgA analysis a Mettler – Toledo TgA/DSC 1 F/500 Stare System apparatus was used, the temperature range was (25 to 600) °C with a heating rate of 10 °C/min under nitrogen flow. All DSC and TgA measurements were performed in triplicate.
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3. Results and discussion 3.1. Normal melting point temperatures and heats of fusion The DSC curves for p-coumaric and caffeic acid, presented in figure 2, show an endothermic peak resulting most probably from melting procedure. The onset temperature determined by DSC method should correspond to normal melting temperature of each acid. However, as shown in the DSC curve of caffeic acid, the endothermic peak is followed by an exothermic one, indicating that another phenomenon except from melting occurs either just after melting is complete or even simultaneously. The TgA curves presented in figure 3 indicate that in both cases acids exhibit a two-stage decomposition, with the first decomposition temperatures at (236 and 241) °C for p-coumaric and caffeic acid, respectively. Comparing the onset temperature obtained by DSC with the decomposition temperature of 1st stage (Table 2), it is concluded that for both acids decomposition takes place following melting, so the temperatures as determined by DSC analysis are actually the normal melting point temperatures of the initial materials. Concerning heat of fusion, although each DSC peak does not correspond only to melting, it would be not a crude approximation to keep those peaks’ integration as a heat of fusion value for each acid. 3.2. Solubility measurements The measured solubilities of two CADs in the six imidazoliumbased ILs and in the two organic solvents in the temperature range of about (303 to 317) K are presented in table 3. Each value represents the average of at least two independent experiments. The corresponding standard deviation (SD) is also reported. Moreover, figure 4 presents the solubilities measured in all solvents studied here at 303 K for a clear picture of the solvent ranking. Despite the solvent variation, caffeic acid is less soluble in all solvents than p-coumaric acid. Caffeic acid’s higher melting temperature and heat of fusion is a reasonable cause. This could also be attributed to the intramolecular hydrogen bonding developed between its two adjacent hydroxyls of the phenyl ring. Since hydroxyls can act as hydrogen bonding donors/acceptors, p-coumaric acid obtains an enhanced energy interaction with a solvent having also hydrogen bonding donors/acceptors. In the case of caf-
feic acid, its two in close distance hydroxyls interact with each other, possibly reducing the interaction from these positions with solvent molecules. As expected due to the very similar structures of the acids, solubility ranking differs slightly; there is only an inverse solubility in omimBF4 and bmimTFO ILs. The solvent ranking for each CAD is as follows: p-Coumaric acid bmimBF4 > bmimTFO > omimBF4 > t-pentanol > ethyl acetate > omimPF6 > bmimPF6 > bmimTF2N. Caffeic acid bmimBF4 > omimBF4 > bmimTFO > t-pentanol > ethyl acetate > omimPF6 > bmimPF6 > bmimTF2N. It is shown that the ILs based on the BF4 and TFO hydrophilic anions are better solvents than those based on PF6 and TF2N hydrophobic anions. One possible reason for this fact is that hydrophilic ILs interact more strongly with the CADs through specific interactions (such as hydrogen bonds). Also the ILs based on BF4 and TFO anions are better solvents than the organic ones, whereas those based on PF6 and TF2N are worse. It is thus proven that with the proper choice of anion and cation the resulting ILs may have higher solvating capacity for CADs than classical organic solvents. Comparing the ILs with the same bmim+ cation, the anion ranking to a decreasing solubility is: BF4 > TFO > PF6 > TF2N. In order to explain this ranking, solvent properties such as polarity and hydrogen bonding must be taken under consideration. There have been made many efforts to quantify ILs’ polarity by various techniques, such as microwave dielectric spectroscopy, chromatographic measurements analyzed using Abraham’s solvation parameter model [31,32], absorption spectra with the use of different solvatochromic dyes (i.e. Nile Red and Reichardt’s dye) [33]. Attempts have also been made to separate non-specific effects of the local electric field from hydrogen-bonding effects using solvation parameters [34]. The literature data have shown that the determination of polarity of ILs is a difficult task and comparison of polarities determined by different techniques that are sensitive to different properties of the solvent often may lead to confusion [35]. Carmichael et al. [33] have shown that the anion polarity ranking for ILs with 1-butyl-3-methyl imidazolium cation is BF4 > TF2-
FIGURE 2. DSC curves for p-coumaric and caffeic acid.
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FIGURE 3. TgA% mass loss and first derivative of % mass loss curves for: (a) p-coumaric acid; (b) caffeic acid.
TABLE 2 Melting temperature, heat of fusion and 1st and 2nd stage decomposition temperatures of p-coumaric (p-CA) and caffeic acid (CA) at pressure p = 0.1 MPa. Compound
Tm/°C ± SD p-CA CA
TGA analysisa
DSC analysis b
219.2 ± 0.3 232.5 ± 0.4
1
DfusH (kJ mol ) ± SD
1st Stage Tdec/°C ± SD
2nd Stage Tdec/°C ± SD
27.42 ± 0.90 27.68 ± 0.13
236.2 ± 0.2 241.1 ± 0.3
388.3 ± 0.4 332.1 ± 0.3
u(T) = ± 0.1, u(DfusH) = ±0.5. Temperature at maximum weight loss determined from first derivative. SD is the standard deviation.
a
b
N > PF6 (polarities probed using the solvatochromic dye Nile Red). However, Chiappe et al. [36] have shown that PF6 anion contributes to higher polarity than TF2N, by using a Michler’s ketone (MK) and tetracyanoethene (TCNE) UV–vis probe so as to investigate donor–acceptor complexes in ILs. To the same result came also Zhang et al. [37], who used the negative solvatochromism of merocyanine as an empirical indicator of the polarity of organic solvents and ionic liquids. Muldoon et al. [38] have measured polarity and nucleophilicity of a range of ionic liquids using two solvatochromic dyes and the results show the following anion polarity ranking for the same bmim+ cation: BF4 > TFO = PF6 > TF2N, whereas bmimPF6 IL is more polar than bmimTFO according to Singh et al. [39]. Taking also into account Kamlet–Taft solvatochromic parameters for bmim+ X ILs (where X = BF4, TFO, PF6, TF2N)
[34,38] it is shown that in the p⁄ scale of dipolarity/polarizability IL ranking is: bmimBF4 > bmimPF6 > bmimTFO > bmimTF2N. In the a-scale of hydrogen bond donor acidity the studied ILs have very similar values, while in the b-scale of hydrogen bond basicity the IL ranking is: bmimTFO > bmimBF4 > bmimTF2N > bmimPF6. It is concluded that although ionic liquids are complex solvents due to their diverse functionalities that lead to diverse interactions with solutes, the IL ranking obtained from the solubility data of the two CADs measured in this work in the bmim+ – based ILs with the four anions is generally in agreement with their relative polarity, hydrophilicity and hydrogen bond basicity. Comparing the ILs with the same anion, it is shown that as the alkyl chain length on the cation increases, the solubility of both CADs decreases in case of ILs with hydrophilic BF4 anion, while increases in case of ILs with hydrophobic PF6 anion. An increase of
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TABLE 3 Mole fraction solubilities, x, of p-coumaric (p-CA) and caffeic acid (CA) in different ILs and in organic solvents at different temperatures and pressure p = 0.1 MPa. Solvent
T/K
p-CA
CA a
x
±SD
x
±SD
bmimPF6
303.1 307.9 312.6 317.4
0.0062 0.0071 0.0084 0.0102
0.00001 0.00004 0.00028 0.00020
0.0012 0.0016 0.0020 0.0024
0.00005 0.00004 0.00001 0.00004
omimPF6
303.1 307.9 312.6 317.4
0.0063 0.0074 0.0090 0.0119
0.00020 0.00001 0.00020 0.00007
0.0014 0.0018 0.0021 0.0024
0.00001 0.00002 0.00007 0.00013
bmimBF4
303.1 307.9 312.6 317.4
0.1951 0.2132 0.2357 0.2513
0.00800 0.00620 0.00580 0.01040
0.1536 0.1637 0.1706 0.1827
0.00330 0.00410 0.00590 0.00110
omimBF4
303.1 307.9 312.6 317.4
0.0803 0.0897 0.0993 0.1100
0.00260 0.00230 0.00370 0.00470
0.0451 0.0518 0.0569 0.0626
0.00010 0.00276 0.00239 0.00162
bmimTF2N
303.1 307.9 312.6 317.4
0.0028 0.0037 0.0042 0.0051
0.00003 0.00013 0.00021 0.00006
0.0003 0.0005 0.0007 0.0011
0.0000039 0.0000044 0.000027 0.000027
bmimTFO
303.1 307.9 312.6 317.4
0.1054 0.1206 0.1333 0.1481
0.00450 0.0040 0.00210 0.00250
0.0242 0.0316 0.0415 0.0541
0.00104 0.00037 0.00093 0.00170
t-Pentanol
303.1 307.9 312.6 317.4
0.0744 0.0774 0.0806 0.0838
0.0003 0.0006 0.0016 0.0013
0.0184 0.0221 0.0264 0.0295
0.0003 0.0001 0.0006 0.0007
Ethyl acetate
303.1 307.9 312.6 317.4
0.0148 0.0164 0.0174 0.0191
0.0004 0.0003 0.0002 0.0004
0.0016 0.0018 0.0021 0.0023
0.00001 0.00002 0.00002 0.00001
u(T) = ±0.1, u(x) = ±0.0005. a SD is the standard deviation.
the hydrophilic parts of CADs and so the solubility decreases, while in the case of hydrophobic ILs this lipophilic change enhances interactions with lipophilic parts of the solute. Finally, solubility of CADs in t-pentanol is higher than in ethyl acetate, which is in accordance with alcohol’s higher polarity and hydrogen bonding capacity. 4. Thermodynamic functions of solution According to van’t Hoff analysis, the apparent standard enthalpy change in solution is obtained from the slope of a ln x2 versus 1/T plot, where x2 is the solute solubility in mole fraction. In recent thermodynamic treatments, some modifications have been introduced to the van’t Hoff treatment to transform the intercept facilitating their use in thermodynamic calculations. According to the Krug et al. [41,42] approach, the following modified van’t Hoff expression is used:
2
3 @ ln x DHoso ln 2 4 5 ¼ : 1 1 R @ TT hm
FIGURE 4. Mole fraction solubility, x, of p-coumaric (p-CA) and caffeic acid (CA) in the studied solvents at 303.05 K.
the alkyl chain length from butyl- to octyl-substitute, invokes a decrease in polarity of the IL [40]. Since CADs’ molecules have both hydrophilic and hydrophobic parts, it is regarded that for hydrophilic ILs, the increased alkyl chain intercepts interactions with
ð1Þ
P
In equation (1), DHosoln/R is the apparent standard enthalpy energy change for the solution process, and Thm is the harmonic mean of the experimental temperatures, which is calculated from:
n T hm ¼ Pn
1 i¼1 T i
;
ð2Þ
where n is the number of temperatures studied. In the present study, Thm is equal to 310.11 K.
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E.I. Alevizou, E.C. Voutsas / J. Chem. Thermodynamics 62 (2013) 69–78 TABLE 4 Apparent thermodynamic functions relative to solution of p-coumaric (p-CA) and caffeic acid (CA) in the studied solvents at pressure p = 0.1 MPa. Solvent
DGosoln/(kJ mol1)
SDa
DHosoln/(kJ mol1)
SD
DSosoln/(J mol1 K1)
bmimPF6 omimPF6 bmimBF4 omimBF4 bmimTF2N bmimTFO t-Pentanol Ethyl acetate
12.50 12.32 3.87 6.09 14.33 5.34 6.55 10.53
0.02 0.03 0.01 0.00 0.03 0.01 0.01 0.01
p-CA 27.9 35.3 14.4 17.6 32.4 18.8 6.7 13.9
0.9 1.8 0.4 0.0 1.6 0.4 0.01 0.4
49.6 74.0 34.1 37.0 58.3 43.5 0.4 10.7
2.9 5.6 1.3 0.1 5.0 1.2 0.1 1.4
bmimPF6 omimPF6 bmimBF4 omimBF4 bmimTF2N bmimTFO t-Pentanol Ethyl acetate
16.38 16.17 4.61 7.54 19.20 8.56 9.65 16.11
0.02 0.02 0.01 0.01 0.03 0.01 0.02 0.01
38.7 29.8 9.4 18.1 71.2 45.1 26.8 20.9
1.2 1.5 0.3 0.5 1.7 0.2 0.9 0.7
72.1 44.0 15.6 34.1 167.6 117.9 55.3 15.4
3.8 4.7 1.0 1.7 5.3 0.8 2.9 2.1
SD
CA
a
SD is the standard deviation.
The application of linear regression models of the modified van’t Hoff plots using the experimentally measured solubilities in this work yielded very high correlation coefficient values (r2 > 0.99). For this reason it is considered that the van’t Hoff equation is useful for estimating the enthalpies of solution in these systems. The apparent standard Gibbs energy change for the solution process (DsolnGo), considering the approach proposed by Krug et al., is calculated at Thm by:
is, the solution process apparently is not spontaneous. Also, the enthalpy of solution is positive in all cases, which suggests that the solution process is always endothermic. Finally, the entropy change is also positive in all cases, indicating that the entropy of solubilization is favorable for both acids in the solvents examined.
Dso ln Go ¼ R T hm K
For the calculation of solid solubility in mole fraction, x2, in a solvent, the following standard thermodynamic equation is applied:
5. Thermodynamic modeling
ð3Þ
in which, the intercept K used is theone obtained in the analysis by treatment of ln x2 as a function of 1T T 1 . hm Finally, the standard apparent entropy change for the solution o process (DsolnS ) is obtained from:
Dso ln So ¼
Dfus H Dfus C p Dfus C p T m T Tm þ ln ðc2 x2 Þ ¼ 1 þ ln 1 Tm R T R T RT ð5Þ
Dso ln Ho Dso ln Go : T hm
ð4Þ where c2, DfusH, and Tm stand for the activity coefficient, the enthalpy of fusion, and the melting temperature of the solid solute, whereas DfusCp is the difference between the heat capacity of the solid and that of the subcooled liquid at the melting temperature.
Table 4 summarizes the apparent standard thermodynamic functions for the solution process of the two CADs in all solvents. The standard Gibbs energy of solution is positive in all cases; that
TABLE 5 Interaction parameters for NRTL and UNIQUAC models for the solubility of CADs (2) in solvent (1). IL
NRTLa (Dg12/R)/K
UNIQUAC (Dg21/R)/K
bmimPF6 omimPF6 bmimBF4 omimBF4 bmimTF2N bmimTFO t-Pentanol Ethyl acetate
195.76 366.92 1301.63 328.89 734.98 964.23 3592.01 2585.55
811.03 94.21 570.11 226.05 198.26 404.46 1025.72 1112.60
bmimPF6 omimPF6 bmimBF4 omimBF4 bmimTF2N bmimTFO t-Pentanol Ethyl acetate
736.31 107.05 276.72 467.76 1063.19 460.68 191.48 229.18
48.76 1183.92 603.84 618.91 58.98 205.44 61.08 3085.36
AAREb p-CA 1.7 4.9 1.4 2.5 2.5 0.7 3.7 7.4
(Du12/R)/K
(Du21/R)/K
AARE
94.91 55.08 357.86 26.22 110.01 259.57 683.59 965.55
185.54 7.08 182.4 131.15 7.20 106.32 382.91 399.18
1.7 5.1 1.3 0.1 2.4 0.9 0.2 2.4
127.13 62.17 12.21 331.36 201.67 138.01 76.51 177.36
4.99 284.08 191.95 437.15 4.43 52.85 24.44 1895.92
3.8 2.7 0.5 0.9 22.0 13.3 1.6 1.0
CA
a b
Nonrandomness of the NRTL model was set equal to 0.2. nP parameter o NP AARE ¼ 1=NP ; NP is the number of points. i¼1 jxexp xcalc j=xexp 100
3.4 2.7 0.5 0.9 21.3 12.8 1.5 2.9
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E.I. Alevizou, E.C. Voutsas / J. Chem. Thermodynamics 62 (2013) 69–78
FIGURE 5. Solubilities of p-coumaric acid in: (a) bmimBF4, bmimTFO, omimBF4 and t-pentanol; (b) bmimPF6, omimPF6, bmimTF2N and ethyl acetate. Solid lines correspond to UNIQUAC calculations.
By neglecting the terms that include DfusCp, equation (5) becomes:
Dfus H T ln ðc2 x2 Þ ¼ 1 Tm RT
ð6Þ
5.1. Correlation with the UNIQUAC & NRTL models Two classical local composition activity coefficient models were used for the correlation of the solubility data, the UNIQUAC [43] and NRTL [44] models. Temperature-independent binary interaction parameters per binary mixture were determined for both models, by minimizing the following objective function (OF):
OF ¼
8 2 91=2 P x2 c2;calc > < NP = 1 > i x c 2 2
> :
NP
> ;
;
ð7Þ
where NP is the number of experimental points, x2 and c2 are the experimental solubility and the corresponding activity coefficient
FIGURE 6. Solubilities of caffeic acid in: (a) bmimBF4, bmimTFO, omimBF4 and tpentanol; (b) bmimPF6, omimPF6, bmimTF2N and ethyl acetate. Solid lines correspond to UNIQUAC calculations.
values, and c2,calc is the activity coefficient calculated with the NRTL or UNIQUAC equation at x = x2,exp. The interaction parameters determined for the two models along with the obtained errors in solubility calculations are presented in table 5, whereas solubility results are graphically presented in figures 5 and 6. It is concluded that very good and similar correlation of the experimental data is obtained by both models, except from the solubilities of caffeic acid in bmimTF2N and bmimTFO that vary strongly with temperature. In these cases temperature-dependent interaction parameters are needed for accurate correlation. Overall, UNIQUAC yields slightly better results than NRTL. Moreover, the effect of the DfusCp, (equation (5)) was examined. Due to the lack of experimental data the value of DfusCp was calculated using predicted C Lp (406 J mol1 K1 for p-coumaric acid and 432 J mol1 K1 for caffeic acid) with the group contribution method of Ruzicka and Domalski [45,46] and predicted C Sp (357 J mol1 K1 for p-coumaric acid and 309 J mol1 K1 for caffeic acid) with the power-law method of Goodman et al. [47]. Unfortunately, worse results were obtained for NRTL and
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E.I. Alevizou, E.C. Voutsas / J. Chem. Thermodynamics 62 (2013) 69–78 TABLE 6 Prediction results with original and modified UNIFAC-Dortmund; mole fraction solubilities, x, of p-coumaric (p-CA) and caffeic acid (CA) in organic solvents. Solvent
T/K
p-CA
CA
x
ARE
a
x
ARE
Original UNIFAC
a
t-Pentanol
303.05 307.85 312.55 317.35
0.0562 0.0609 0.0658 0.0712
24.4 21.3 18.3 15.1
0.0886 0.0927 0.0969 0.1014
381.4 319.3 266.9 243.7
Ethyl acetate
303.05 307.85 312.55 317.35
0.0561 0.0618 0.0678 0.0744
279.1 276.8 289.7 289.5
0.0863 0.0915 0.0969 0.1027
5293.1 4981.7 4512.4 4365.2
t-Pentanol
303.05 307.85 312.55 317.35
Modified UNIFAC-Dortmund 0.0124 0.0158 0.0199 0.0249
83.3 79.6 75.3 70.3
0.0142 0.0182 0.0228 0.0283
22.9 17.8 13.7 4.2
Ethyl acetate
303.05 307.85 312.55 317.35
0.0005 0.0006 0.0008 0.0010
96.6 96.1 95.4 94.6
0.0694 0.0753 0.0813 0.0879
4240.0 4082.2 3773.3 3721.3
ARE ¼ jxexp xcalc j=xexp 100.
UNIQUAC when the DfusCp terms were taken into account that may be due to the predicted DfusCp values used.
5.2. Prediction of CADs solubility in Organic Solvents with UNIFAC The original UNIFAC model [48] and the modified UNIFAC-Dortmund [49] were applied for the prediction of CADs solubility only in tert-pentanol and ethyl-acetate, since UNIFAC parameters are missing for ILs. The results are presented in table 6. Original UNIFAC erroneously predicts higher solubilities of p-coumaric acid in ethyl acetate than in t-pentanol, except from the lowest temperature, while it correctly predicts higher solubilities of caffeic acid in t-pentanol than in ethyl acetate, except from the highest temperature. On the other hand, modified UNIFAC-Dortmund correctly predicts the ranking of the two solvents for the case of p-coumaric acid, while it predicts the opposite trend for caffeic acid. Nevertheless, both models give in most cases very poor qualitative predictions.
6. Conclusions The solubilities of p-coumaric acid and caffeic acid have been measured in six alkyl-substituted imidazolium-based ILs composed of the BF4, PF6, TF2N and TFO anions and in two organic solvents; t-pentanol and ethyl acetate. The results showed that the BF4 and TFO – based ILs are better solvents than the organic ones, which in turn are better than the corresponding PF6 and TF2N – based ILs. Furthermore, it is noticed that the increase of the alkyl chain length on the cation leads to a decrease in acids’ solubility in case of hydrophilic ILs with BF4 anion, but to an increase in case of hydrophobic ILs with PF6 anion. Moreover, melting temperatures and heats of fusion of the two acids were determined by differential scanning calorimetry, while thermogravimetric analysis revealed that both acids decompose forthwith melting. By utilizing the experimental solubilities, the thermodynamic functions of solutions were determined. The results indicated that the solution process is not spontaneous, is always endothermic and the entropy of solubilization is favorable for all systems studied. Finally, successful correlation with the NRTL and UNIQUAC models was obtained for all systems studied, with UNIQUAC being overall slightly
better than NRTL, while poor predictions of the two acids’ solubility in the organic solvents were obtained with two UNIFAC models. Acknowledgements The financial support of the National Technical University of Athens, Greece (program ‘‘PEVE 2009’’ for basic research) is gratefully acknowledged. References [1] M.N. Clifford, J. Sci. Food Agric. 79 (1999) 362–372. [2] Y. Cui, Q. Li, M. Zhang, Z. Liu, W. Yin, W. Liu, X. Chen, K. Bi, J. Agric. Food Chem. 58 (2010) 12083–12088. [3] S. Vishnoi, V. Agrawal, V.K. Kasana, J. Agric. Food Chem. 57 (2009) 3261–3265. [4] F.A. Ayaz, S. Hayirlioglu-Ayaz, J. Gruz, O. Novak, M. Strnad, J. Agric. Food Chem. 53 (2005) 8116–8122. [5] J.L. Slavin, J. Am. Coll. Nutr. 19 (2000) 300s–307s. [6] I.C.W. Arts, P.C.H. Hollman, Am. J. Clin. Nutr. 81 (2005) 317s–325s. [7] C.A. Rice-Evans, N.J. Miller, G. Paganga, Free Radical Biol. Med. 20 (1996) 933– 956. [8] F.t.L. Mota, A.n.J. Queimada, S.o.P. Pinho, E.n.A. Macedo, Ind. Eng. Chem. Res. 47 (2008) 5182–5189. [9] C. Fabian, N.Y. Tran-Thi, N.S. Kasim, Y.H. Ju, J. Sci. Food Agric. 90 (2010) 2576– 2581. [10] A. Apelblat, E. Manzurola, N.A. Balal, J. Chem. Thermodyn. 38 (2006) 565–571. [11] C.M.S.S. Neves, M.L.S. Batista, A.F.M. Claudio, L.M.N.B.F. Santos, I.M. Marrucho, M.G. Freire, M.G. Freire, J.A.P. Coutinho, J. Chem. Eng. Data. 55 (2010) 5065– 5073. [12] M.E. Zakrzewska, E. Bogel-Łukasik, R. Bogel-Łukasik, Energy Fuels 24 (2010) 737–745. [13] P. Migowski, J. Dupont, Chem. – A Eur. J. 13 (2007) 32–39. [14] K.N. Marsh, J.A. Boxall, R. Lichtenthaler, Fluid Phase Equilib. 219 (2004) 93–98. [15] L. Chebil, C. Humeau, J. Anthoni, F. Dehez, J.-M. Engasser, M. Ghoul, J. Chem. Eng. Data 52 (2007) 1552–1556. [16] B. Peng, J. Zi, W. Yan, J. Chem. Eng. Data 51 (2006) 2038–2040. [17] A.J. Queimada, F.L. Mota, S.P. Pinho, E.A. Macedo, J. Phys. Chem. B 113 (2009) 3469–3476. [18] W. He, D.S. Silvester, I. Streeter, L. Aldous, C. Hardacre, R.G. Compton, J. Phys. Org. Chem. 22 (2009) 69–76. [19] E.K. Panteli, E.K. Voutsas, J. Chem. Eng. Data 54 (2009) 812–818. [20] R. Bogel-Lukasik, L.M. Nobre Goncalves, E. Bogel-Lukasik, Green Chem. 12 (2010) 1947–1953. [21] M.S. Manic, V. Najdanovic-Visak, J. Chem. Thermodyn. 44 (2012) 102–106. [22] H.-M. Lin, H.-Y. Tien, Y.-T. Hone, M.-J. Lee, Fluid Phase Equilib. 253 (2007) 130– 136. [23] C. Lourenço, C.I. Melo, R. Bogel-Łukasik, E. Bogel-Łukasik, J. Chem. Eng. Data 57 (2012) 1525–1533. [24] A. Forte, C.I. Melo, R. Bogel-Łukasik, E. Bogel-Łukasik, Fluid Phase Equilib. 318 (2012) 89–95. [25] K.B. Smith, R.H. Bridson, G.A. Leeke, J. Chem. Eng. Data 56 (2011) 2039–2043.
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JCT 12-431
Contents lists available at SciVerse ScienceDirect
J. Chem. Thermodynamics journal homepage: www.elsevier.com/locate/jct
Solubilities of p-coumaric and caffeic acid in ionic liquids and organic solvents Efthimia I. Alevizou, Epaminondas C. Voutsas ⇑ Laboratory of Thermodynamics and Transport Phenomena, School of Chemical Engineering, National Technical University of Athens, 9 Heroon Polytechniou Str., Zographou Campus, 15780 Athens, Greece
a r t i c l e
i n f o
Article history: Received 24 July 2012 Received in revised form 21 February 2013 Accepted 23 February 2013 Available online 4 March 2013 Keywords: Antioxidants p-Coumaric acid Caffeic acid Ionic liquids Solubility measurements Correlation
a b s t r a c t The solubilities of two cinnamic acid derivatives, namely p-coumaric acid and caffeic acid, in six 1alkyl-3-methyl imidazolium based ionic liquids composed of the PF6, BF4, TFO and TF2N anions, and in two organic solvents, t-pentanol and ethyl acetate, have been measured at the temperature range of about (303 to 317) K. The p-coumaric acid was found to be more soluble than caffeic acid in all studied solvents. Higher solubilities of both acids were observed in the ionic liquids composed of the BF4 and TFO anions. The increase of the alkyl chain length on the cation invokes a decrease in solubility in the case of hydrophilic ionic liquids composed of BF4 anion, while in the case of hydrophobic ones composed of PF6 anion an increase in the solubility is observed. Between the two organic solvents t-pentanol is better solvent than ethyl acetate for both acids. Moreover, using the van’t Hoff equations the apparent Gibbs energy, enthalpy, and entropy of solution were calculated. Finally, successful correlation of the experimental data was achieved with the UNIQUAC and the NRTL activity coefficient models, while poor predictions of the solubility of the two acids in the organic solvents were obtained with two UNIFAC models. Ó 2013 Elsevier Ltd. All rights reserved.
1. Introduction Cinnamic acid derivatives (CADs), such as p-coumaric acid and caffeic acid, are natural antioxidants belonging to the family of phenolic acids and are widely distributed throughout the plant kingdom. Fruits, vegetables, spices, aromatic herbs, cereals, coffee beans and beverages, olives, propolis, sunflower, barks and wine are natural sources of these compounds [1–4]. Due to their almost universal distribution, they are an integral part of the human diet. Epidemiological studies suggest a link between the consumption of whole grain products containing CADs and prevention of chronic, degenerative diseases associated with oxidative damage (namely coronary heart diseases, cardiovascular diseases, diabetes, arthritis, cataract formation, aging and cancer) [5,6]. Many of the health protective effects of phenolic compounds have been ascribed to their antioxidant, antimutagenic, anti-inflammatory, antimicrobial, antiviral, antialergic, immunoprotective, ultraviolet (UV) filtering properties and other biological and pharmacological activities Abbreviations: CADs, cinnamic acid derivatives; ILs, ionic liquids; p-CA, p-coumaric acid; CA, caffeic acid; bmim, 1-butyl-3-methyl imidazolium; omim, 1-octyl-3-methyl imidazolium; PF6, hexafluorophosphate; BF4, tetrafluoroborate; TF2N, bis(trifluoromethanesulfonyl)imide; TFO, trifluromethanesulfonate; HPLC, high-performance liquid chromatography; DSC, differential scanning calorimetry; TgA, thermogravimetric analysis; calc, calculated; exp, experimental; SD, standard deviation; AARE, absolute average relative error; ARE, absolute relative error. ⇑ Corresponding author. Tel.: +30 210 772 3971; fax: +30 210 772 3155. E-mail address: [email protected] (E.C. Voutsas). 0021-9614/$ - see front matter Ó 2013 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.jct.2013.02.013
[7,8]. The above benefits in combination with a growing concern about the safety of synthetic antioxidants [9] make CADs high value raw materials for the synthesis of different molecules with industrial interest, such as drugs, cosmetics, antiseptics and flavors [10]. Common problems encountered in chemical synthesis, recovery and separation processes, such as the use of toxic, volatile and flammable organic solvents as well as difficulties in product purification can be handled with the use of ionic liquids (ILs); a new class of solvents that can lead to more environmentally friendly applications and processes as compared to those where conventional organic solvents are used. This characteristic arises mainly from their negligible vapor pressures at room temperature and therefore low volatility and flammability, which minimizes health and safety risk in industry and the chance of loss to atmosphere making their recycling and reusability feasible. Of course, the development of cost and energy efficient technologies for solute recovery from ionic liquids is a challenge towards their industrial application. Ionic liquids are molten salts, consisting of large asymmetric organic cations and organic or inorganic anions. The low symmetry, high vibrational freedom and charge delocalization of the ions composing an IL reduce the stability of the crystalline phases, and thus their melting temperatures [11]. Unlike molecular liquids, their ionic nature results in a unique combination of distinctive properties, such as high thermal stabilities, large liquidus range and high solvating capacity for organic, inorganic and
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organometallic compounds [12,13]. The last is one of the most attractive features of ILs especially for recovery and separation processes, as well as final product purification. It results from ILs’ ability to be tailormade for a specific purpose by careful selection among a huge diversity of cations and anions, or substitutes to the cation. Therefore, ILs are often referred to as ‘‘designer solvents’’ [14]. For the design and optimization of processes where CADs are involved, suitable solvents – either classical organic ones or ionic liquids – have to be selected. The successful completion of such a task requires the knowledge of reliable phase equilibrium data. With reference to solubilities of antioxidants in organic solvents, there are some data in the literature, such as solubility of flavonoids in acetonitrile, acetone and tert-amyl alcohol [15], or solubility of luteolin in methanol, ethanol, 1-propanol, 2-propanol, 1-butanol, acetone, hexane, and DMSO [16], but for phenolic acids only their solubility in water has been measured [8,17]. Concerning ILs, although much experimental work has been conducted so far on measurements of thermophysical properties of the pure ILs and phase equilibrium in systems containing ILs, solubility data of antioxidants in ILs are very scarce [18–20]. Some other works have also reported solubility data in ILs of antibiotics [21], dibasic carboxylic acids [22], drugs [23–25] and inorganic salts [26]. In this work, the solubilities of two CADs, namely p-coumaric acid and caffeic acid, were measured in six imidazolium based ILs composed of the PF6, BF4, TF2N and TFO anions and they were compared to those measured in two organic solvents, t-pentanol and ethyl acetate. In addition, the normal melting point temperature and the enthalpy of fusion of the two acids were determined by differential scanning calorimentry and their possible decomposition upon melting was tested by thermogravimetric analysis. Moreover, by utilizing the solubility data and the van’t Hoff equations, the apparent thermodynamic functions relative to the solution of the two CADs have been determined. Finally, the capability of the UNIQUAC and the NRTL activity coefficient models to correlate the solubility data was tested and two UNIFAC versions were used to predict the solubilities of the two acids in the organic solvents.
2. Experimental section 2.1. Materials p-Coumaric acid (p-CA; purity > 98%; CAS No. 501-98-4; C9H8O3) was purchased from Sigma Aldrich Co. Caffeic acid (CA, purity > 98% CAS No. 331-39-5; C9H8O4) was purchased from Acros
Organics. 2-methyl-2-butanol (t-pentanol; purity > 99%; CAS No. 75-85-4; C5H12O) was purchased from Sigma Aldrich Co. Ethyl acetate (purity > 99.8%; CAS No. 141-78-6; C4H8O2) was purchased from Merck KGaA. 1-butyl-3-methyl-imidazolium hexafluorophosphate (bmimPF6; purity > 98%; CAS No. 174501-64-5; C8H15F6N2P), 1-octyl-3-methyl-imidazolium hexafluorophosphate (omimPF6; purity > 98%; CAS No. 304680-36-2; C12H23F6PN2), 1-butyl-3methyl-imidazolium tetrafluoroborate (bmimBF4; purity > 98%, CAS No. 174501-65-6; C8H15BF4N2), 1-octyl-3-methyl-imidazolium tetrafluoroborate (omimBF4; purity > 98%; CAS No. 244193-52-0; C12H23BF4N2), 1-butyl-3-methyl-imidazolium bis(trifluoromethanesulfonyl)imide (bmimTF2N; purity > 98%, CAS No. 17489983-3; C10H15F6N3O4S2), 1-butyl-3-methyl-imidazolium trifluoromethanesulfonate (bmimTFO, purity > 98%, CAS No. 174899-66-2; C9H15F3N2O3S) were purchased from Solchemar Lda., Portugal. Methanol (99.9% purity) and acetonitrile (99.9% purity) were purchased from Fisher Scientific. Water (HPLC gradient) was purchased from Merck. Table 1 summarizes the source and purities of compounds used in this work. The molecular structures of the investigated cinnamic acid derivatives (CADs) and the cations and anions of the ionic liquids (ILs) are presented in figure 1. 2.2. Solubility measurements and procedure The equilibrium runs were performed in a Thermomixer Comfort (Eppendorf). This apparatus combines stirring up to 1400 rpm and heating (up to 99 °C) or cooling (13 °C below room temperature). The temperature readings of the thermomixer have been checked using an external thermometer calibrated with the Isocal-6 VenusPlus 2140 calibrator of temperature sensors. The accompanied thermoblock can accommodate 24 micro test tubes of 2 mL each. Approximately 200 lL of each solvent were placed in the 2 mL volume safe-lock micro test tube. An excess quantity of each acid was then added. The tubes were placed in the apparatus, which was set to the desired temperature measured with an uncertainty of ±0.1 K, and were stirred at 1400 rpm until equilibrium was reached. Based on Mota et al.’s solubility measurements of caffeic acid in water [8], where a quite long shaking time was needed (64 to 84) h, it was concluded that in our case of indirect stirring and more viscous solvents the equilibrium times would be even longer. To determine the time required to reach equilibrium, a preliminary measurement of the solubility of the most soluble acid (p-CA) in the most viscous IL omimBF4 was performed. Samples were taken and analyzed after (24, 72, 120, 144, 168, and 192) h of stirring. It was found that the difference among the solubility values measured after (144, 168, and 192) h was less than 2%,
TABLE 1 List of compounds, their source and purities.
a b
Chemical name
Source
Purity %
3-(4-Hydroxyphenyl)-2-propenoic acid a 3-(3,4-Dihydroxyphenyl)-2-propenoic acid b 1-Butyl-3-methyl-imidazolium hexafluorophosphate 1-Octyl-3-methyl-imidazolium hexafluorophosphate 1-Butyl-3-methyl-imidazolium tetrafluoroborate 1-Octyl-3-methyl-imidazolium tetrafluoroborate 1-Butyl-3-methyl-imidazolium bis(trifluoromethanesulfonyl)imide 1-Butyl-3-methyl-imidazolium trifluoromethanesulfonate 2-Methyl-2-butanol Ethyl acetate Methanol Acetonitrile Water
Sigma Aldrich Co Acros Organics Solchemar Lda Solchemar Lda Solchemar Lda Solchemar Lda Solchemar Lda Solchemar Lda Sigma Aldrich Co Merck KGaA Fisher Scientific Fisher Scientific Merck KGaA
98.0 98.0 98.0 98.0 98.0 98.0 98.0 98.0 99.0 99.8 99.9 99.9 99.99
p-Coumaric acid. Caffeic acid.
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OH
OH OH
OH
OH
O
p-coumaric acid (p-CA)
O
caffeic acid (CA) +
+
N
N
N
N
1-butyl-3-methyl imidazolium cation (bmim+) 1-octyl-3-methyl imidazolium cation (omim+)
hexafluorophosphate anion (PF6-)
bis(trifluoromethylsulfonyl)imide anion (TF2N-)
tetrafluoroborate anion (BF4-)
trifluoromethylsulfonate anion (TFO-)
FIGURE 1. Structures of investigated cinnamic acid derivatives and cations/anions of ionic liquids.
which led to the conclusion that a time frame of 144 h is sufficient for achieving equilibrium for all solutions. After equilibrium was reached, solution samples were withdrawn using adjustable micropipettes. The sampling tips of the pipettes were preheated at the temperature of the experiment to avoid solid precipitation. Before sampling, each solution was centrifuged for a short time (60 to 90) s to enhance the physical separation of the two phases, liquid and solid, and to ensure that no precipitated solute was also drawn away during sampling because of the high viscosities of the ILs. All experiments were performed in triplicate. The water content in the studied solvents was measured by Karl-Fisher titration (volumetric titrator, TitroLine KF, SCHOTT Instruments) and was found to be: 0.97% w/w for bmimBF4, 1.45% for omimBF4, 1.35% for bmimPF6, 1.4% for omimPF6, 0.98% for bmimTF2N, 1.55% for bmimTFO, 0.06% for t-pentanol and 0.14% for ethyl acetate. The effect of the water traces in the ILs on the solubility was also investigated. For this purpose, solubility measurements were performed using dried (at 80 °C under vacuum for 48 h) ILs, and when they were compared with those obtained in the ILs used as purchased, no difference was found. For solubility measurements all solvents were used as received. Samples were analyzed by reverse phase high-performance liquid chromatography (HPLC). The HPLC system was equipped with a Jusco PU-1580 pump, an ICI LC 1200 UV/vis detector set at 320 nm, and a BDS Hypersil C18 column ((250 4.6) mm2, particle size 5 lm, Thermo Scientific). The eluent solution used was a mixture of acetonitrile/water (80/20 v/v) with 0.1% acetic acid. The flow rate was adjusted to 1 mL min1, the column temperature was kept constant at 30 °C, and the injection volume was 50 lL. The eluent solution was degassed with Helium.
2.3. Differential scanning calorimetry (DSC) and thermogravimetric analysis (TgA) For the calculation of the solid solubility, the normal melting point temperature (Tm) and heat of fusion (DfusH) of p-coumaric and caffeic acid are needed. Although there are such data in the literature, they show notable differences mostly regarding the melting point values and whether or not the two acids decompose upon melting. For example, Park et al. [27] have determined the melting point temperature of caffeic acid (Tm = 232.32 °C) with a DSC analysis without referring to a possible decomposition, while Mota et al. [8] could not determine it – also with a DSC analysis – because of its decomposition upon melting. The NIST Chemistry Webbook [28] reports a melting temperature range of (202 to 216) °C for p-coumaric acid. Saldana et al. [29] present the Tm of p-coumaric acid to be 214 °C and that of caffeic acid to be 196 °C, both taken from the Material Safety Data Sheets of Alfa and Sigma–Aldrich, respectively. The same values are presented also by Murga et al. [30] taken from Sigma–Aldrich. It is therefore concluded that there is a high uncertainty for the normal melting point temperatures of the two CADs. Thus, it was necessary to carry out a DSC analysis in combination with a TgA one, so as to determine the Tm and DfusH of the studied materials and also to investigate possible decomposition upon melting. For DSC analysis a Mettler – Toledo DSC1 apparatus was used, the temperature range was (50 to 300) °C with a heating rate of 10 °C/min under nitrogen flow. DSC apparatus was calibrated with indium, zinc and n-heptane at the onset temperature. For TgA analysis a Mettler – Toledo TgA/DSC 1 F/500 Stare System apparatus was used, the temperature range was (25 to 600) °C with a heating rate of 10 °C/min under nitrogen flow. All DSC and TgA measurements were performed in triplicate.
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3. Results and discussion 3.1. Normal melting point temperatures and heats of fusion The DSC curves for p-coumaric and caffeic acid, presented in figure 2, show an endothermic peak resulting most probably from melting procedure. The onset temperature determined by DSC method should correspond to normal melting temperature of each acid. However, as shown in the DSC curve of caffeic acid, the endothermic peak is followed by an exothermic one, indicating that another phenomenon except from melting occurs either just after melting is complete or even simultaneously. The TgA curves presented in figure 3 indicate that in both cases acids exhibit a two-stage decomposition, with the first decomposition temperatures at (236 and 241) °C for p-coumaric and caffeic acid, respectively. Comparing the onset temperature obtained by DSC with the decomposition temperature of 1st stage (Table 2), it is concluded that for both acids decomposition takes place following melting, so the temperatures as determined by DSC analysis are actually the normal melting point temperatures of the initial materials. Concerning heat of fusion, although each DSC peak does not correspond only to melting, it would be not a crude approximation to keep those peaks’ integration as a heat of fusion value for each acid. 3.2. Solubility measurements The measured solubilities of two CADs in the six imidazoliumbased ILs and in the two organic solvents in the temperature range of about (303 to 317) K are presented in table 3. Each value represents the average of at least two independent experiments. The corresponding standard deviation (SD) is also reported. Moreover, figure 4 presents the solubilities measured in all solvents studied here at 303 K for a clear picture of the solvent ranking. Despite the solvent variation, caffeic acid is less soluble in all solvents than p-coumaric acid. Caffeic acid’s higher melting temperature and heat of fusion is a reasonable cause. This could also be attributed to the intramolecular hydrogen bonding developed between its two adjacent hydroxyls of the phenyl ring. Since hydroxyls can act as hydrogen bonding donors/acceptors, p-coumaric acid obtains an enhanced energy interaction with a solvent having also hydrogen bonding donors/acceptors. In the case of caf-
feic acid, its two in close distance hydroxyls interact with each other, possibly reducing the interaction from these positions with solvent molecules. As expected due to the very similar structures of the acids, solubility ranking differs slightly; there is only an inverse solubility in omimBF4 and bmimTFO ILs. The solvent ranking for each CAD is as follows: p-Coumaric acid bmimBF4 > bmimTFO > omimBF4 > t-pentanol > ethyl acetate > omimPF6 > bmimPF6 > bmimTF2N. Caffeic acid bmimBF4 > omimBF4 > bmimTFO > t-pentanol > ethyl acetate > omimPF6 > bmimPF6 > bmimTF2N. It is shown that the ILs based on the BF4 and TFO hydrophilic anions are better solvents than those based on PF6 and TF2N hydrophobic anions. One possible reason for this fact is that hydrophilic ILs interact more strongly with the CADs through specific interactions (such as hydrogen bonds). Also the ILs based on BF4 and TFO anions are better solvents than the organic ones, whereas those based on PF6 and TF2N are worse. It is thus proven that with the proper choice of anion and cation the resulting ILs may have higher solvating capacity for CADs than classical organic solvents. Comparing the ILs with the same bmim+ cation, the anion ranking to a decreasing solubility is: BF4 > TFO > PF6 > TF2N. In order to explain this ranking, solvent properties such as polarity and hydrogen bonding must be taken under consideration. There have been made many efforts to quantify ILs’ polarity by various techniques, such as microwave dielectric spectroscopy, chromatographic measurements analyzed using Abraham’s solvation parameter model [31,32], absorption spectra with the use of different solvatochromic dyes (i.e. Nile Red and Reichardt’s dye) [33]. Attempts have also been made to separate non-specific effects of the local electric field from hydrogen-bonding effects using solvation parameters [34]. The literature data have shown that the determination of polarity of ILs is a difficult task and comparison of polarities determined by different techniques that are sensitive to different properties of the solvent often may lead to confusion [35]. Carmichael et al. [33] have shown that the anion polarity ranking for ILs with 1-butyl-3-methyl imidazolium cation is BF4 > TF2-
FIGURE 2. DSC curves for p-coumaric and caffeic acid.
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FIGURE 3. TgA% mass loss and first derivative of % mass loss curves for: (a) p-coumaric acid; (b) caffeic acid.
TABLE 2 Melting temperature, heat of fusion and 1st and 2nd stage decomposition temperatures of p-coumaric (p-CA) and caffeic acid (CA) at pressure p = 0.1 MPa. Compound
Tm/°C ± SD p-CA CA
TGA analysisa
DSC analysis b
219.2 ± 0.3 232.5 ± 0.4
1
DfusH (kJ mol ) ± SD
1st Stage Tdec/°C ± SD
2nd Stage Tdec/°C ± SD
27.42 ± 0.90 27.68 ± 0.13
236.2 ± 0.2 241.1 ± 0.3
388.3 ± 0.4 332.1 ± 0.3
u(T) = ± 0.1, u(DfusH) = ±0.5. Temperature at maximum weight loss determined from first derivative. SD is the standard deviation.
a
b
N > PF6 (polarities probed using the solvatochromic dye Nile Red). However, Chiappe et al. [36] have shown that PF6 anion contributes to higher polarity than TF2N, by using a Michler’s ketone (MK) and tetracyanoethene (TCNE) UV–vis probe so as to investigate donor–acceptor complexes in ILs. To the same result came also Zhang et al. [37], who used the negative solvatochromism of merocyanine as an empirical indicator of the polarity of organic solvents and ionic liquids. Muldoon et al. [38] have measured polarity and nucleophilicity of a range of ionic liquids using two solvatochromic dyes and the results show the following anion polarity ranking for the same bmim+ cation: BF4 > TFO = PF6 > TF2N, whereas bmimPF6 IL is more polar than bmimTFO according to Singh et al. [39]. Taking also into account Kamlet–Taft solvatochromic parameters for bmim+ X ILs (where X = BF4, TFO, PF6, TF2N)
[34,38] it is shown that in the p⁄ scale of dipolarity/polarizability IL ranking is: bmimBF4 > bmimPF6 > bmimTFO > bmimTF2N. In the a-scale of hydrogen bond donor acidity the studied ILs have very similar values, while in the b-scale of hydrogen bond basicity the IL ranking is: bmimTFO > bmimBF4 > bmimTF2N > bmimPF6. It is concluded that although ionic liquids are complex solvents due to their diverse functionalities that lead to diverse interactions with solutes, the IL ranking obtained from the solubility data of the two CADs measured in this work in the bmim+ – based ILs with the four anions is generally in agreement with their relative polarity, hydrophilicity and hydrogen bond basicity. Comparing the ILs with the same anion, it is shown that as the alkyl chain length on the cation increases, the solubility of both CADs decreases in case of ILs with hydrophilic BF4 anion, while increases in case of ILs with hydrophobic PF6 anion. An increase of
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TABLE 3 Mole fraction solubilities, x, of p-coumaric (p-CA) and caffeic acid (CA) in different ILs and in organic solvents at different temperatures and pressure p = 0.1 MPa. Solvent
T/K
p-CA
CA a
x
±SD
x
±SD
bmimPF6
303.1 307.9 312.6 317.4
0.0062 0.0071 0.0084 0.0102
0.00001 0.00004 0.00028 0.00020
0.0012 0.0016 0.0020 0.0024
0.00005 0.00004 0.00001 0.00004
omimPF6
303.1 307.9 312.6 317.4
0.0063 0.0074 0.0090 0.0119
0.00020 0.00001 0.00020 0.00007
0.0014 0.0018 0.0021 0.0024
0.00001 0.00002 0.00007 0.00013
bmimBF4
303.1 307.9 312.6 317.4
0.1951 0.2132 0.2357 0.2513
0.00800 0.00620 0.00580 0.01040
0.1536 0.1637 0.1706 0.1827
0.00330 0.00410 0.00590 0.00110
omimBF4
303.1 307.9 312.6 317.4
0.0803 0.0897 0.0993 0.1100
0.00260 0.00230 0.00370 0.00470
0.0451 0.0518 0.0569 0.0626
0.00010 0.00276 0.00239 0.00162
bmimTF2N
303.1 307.9 312.6 317.4
0.0028 0.0037 0.0042 0.0051
0.00003 0.00013 0.00021 0.00006
0.0003 0.0005 0.0007 0.0011
0.0000039 0.0000044 0.000027 0.000027
bmimTFO
303.1 307.9 312.6 317.4
0.1054 0.1206 0.1333 0.1481
0.00450 0.0040 0.00210 0.00250
0.0242 0.0316 0.0415 0.0541
0.00104 0.00037 0.00093 0.00170
t-Pentanol
303.1 307.9 312.6 317.4
0.0744 0.0774 0.0806 0.0838
0.0003 0.0006 0.0016 0.0013
0.0184 0.0221 0.0264 0.0295
0.0003 0.0001 0.0006 0.0007
Ethyl acetate
303.1 307.9 312.6 317.4
0.0148 0.0164 0.0174 0.0191
0.0004 0.0003 0.0002 0.0004
0.0016 0.0018 0.0021 0.0023
0.00001 0.00002 0.00002 0.00001
u(T) = ±0.1, u(x) = ±0.0005. a SD is the standard deviation.
the hydrophilic parts of CADs and so the solubility decreases, while in the case of hydrophobic ILs this lipophilic change enhances interactions with lipophilic parts of the solute. Finally, solubility of CADs in t-pentanol is higher than in ethyl acetate, which is in accordance with alcohol’s higher polarity and hydrogen bonding capacity. 4. Thermodynamic functions of solution According to van’t Hoff analysis, the apparent standard enthalpy change in solution is obtained from the slope of a ln x2 versus 1/T plot, where x2 is the solute solubility in mole fraction. In recent thermodynamic treatments, some modifications have been introduced to the van’t Hoff treatment to transform the intercept facilitating their use in thermodynamic calculations. According to the Krug et al. [41,42] approach, the following modified van’t Hoff expression is used:
2
3 @ ln x DHoso ln 2 4 5 ¼ : 1 1 R @ TT hm
FIGURE 4. Mole fraction solubility, x, of p-coumaric (p-CA) and caffeic acid (CA) in the studied solvents at 303.05 K.
the alkyl chain length from butyl- to octyl-substitute, invokes a decrease in polarity of the IL [40]. Since CADs’ molecules have both hydrophilic and hydrophobic parts, it is regarded that for hydrophilic ILs, the increased alkyl chain intercepts interactions with
ð1Þ
P
In equation (1), DHosoln/R is the apparent standard enthalpy energy change for the solution process, and Thm is the harmonic mean of the experimental temperatures, which is calculated from:
n T hm ¼ Pn
1 i¼1 T i
;
ð2Þ
where n is the number of temperatures studied. In the present study, Thm is equal to 310.11 K.
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E.I. Alevizou, E.C. Voutsas / J. Chem. Thermodynamics 62 (2013) 69–78 TABLE 4 Apparent thermodynamic functions relative to solution of p-coumaric (p-CA) and caffeic acid (CA) in the studied solvents at pressure p = 0.1 MPa. Solvent
DGosoln/(kJ mol1)
SDa
DHosoln/(kJ mol1)
SD
DSosoln/(J mol1 K1)
bmimPF6 omimPF6 bmimBF4 omimBF4 bmimTF2N bmimTFO t-Pentanol Ethyl acetate
12.50 12.32 3.87 6.09 14.33 5.34 6.55 10.53
0.02 0.03 0.01 0.00 0.03 0.01 0.01 0.01
p-CA 27.9 35.3 14.4 17.6 32.4 18.8 6.7 13.9
0.9 1.8 0.4 0.0 1.6 0.4 0.01 0.4
49.6 74.0 34.1 37.0 58.3 43.5 0.4 10.7
2.9 5.6 1.3 0.1 5.0 1.2 0.1 1.4
bmimPF6 omimPF6 bmimBF4 omimBF4 bmimTF2N bmimTFO t-Pentanol Ethyl acetate
16.38 16.17 4.61 7.54 19.20 8.56 9.65 16.11
0.02 0.02 0.01 0.01 0.03 0.01 0.02 0.01
38.7 29.8 9.4 18.1 71.2 45.1 26.8 20.9
1.2 1.5 0.3 0.5 1.7 0.2 0.9 0.7
72.1 44.0 15.6 34.1 167.6 117.9 55.3 15.4
3.8 4.7 1.0 1.7 5.3 0.8 2.9 2.1
SD
CA
a
SD is the standard deviation.
The application of linear regression models of the modified van’t Hoff plots using the experimentally measured solubilities in this work yielded very high correlation coefficient values (r2 > 0.99). For this reason it is considered that the van’t Hoff equation is useful for estimating the enthalpies of solution in these systems. The apparent standard Gibbs energy change for the solution process (DsolnGo), considering the approach proposed by Krug et al., is calculated at Thm by:
is, the solution process apparently is not spontaneous. Also, the enthalpy of solution is positive in all cases, which suggests that the solution process is always endothermic. Finally, the entropy change is also positive in all cases, indicating that the entropy of solubilization is favorable for both acids in the solvents examined.
Dso ln Go ¼ R T hm K
For the calculation of solid solubility in mole fraction, x2, in a solvent, the following standard thermodynamic equation is applied:
5. Thermodynamic modeling
ð3Þ
in which, the intercept K used is theone obtained in the analysis by treatment of ln x2 as a function of 1T T 1 . hm Finally, the standard apparent entropy change for the solution o process (DsolnS ) is obtained from:
Dso ln So ¼
Dfus H Dfus C p Dfus C p T m T Tm þ ln ðc2 x2 Þ ¼ 1 þ ln 1 Tm R T R T RT ð5Þ
Dso ln Ho Dso ln Go : T hm
ð4Þ where c2, DfusH, and Tm stand for the activity coefficient, the enthalpy of fusion, and the melting temperature of the solid solute, whereas DfusCp is the difference between the heat capacity of the solid and that of the subcooled liquid at the melting temperature.
Table 4 summarizes the apparent standard thermodynamic functions for the solution process of the two CADs in all solvents. The standard Gibbs energy of solution is positive in all cases; that
TABLE 5 Interaction parameters for NRTL and UNIQUAC models for the solubility of CADs (2) in solvent (1). IL
NRTLa (Dg12/R)/K
UNIQUAC (Dg21/R)/K
bmimPF6 omimPF6 bmimBF4 omimBF4 bmimTF2N bmimTFO t-Pentanol Ethyl acetate
195.76 366.92 1301.63 328.89 734.98 964.23 3592.01 2585.55
811.03 94.21 570.11 226.05 198.26 404.46 1025.72 1112.60
bmimPF6 omimPF6 bmimBF4 omimBF4 bmimTF2N bmimTFO t-Pentanol Ethyl acetate
736.31 107.05 276.72 467.76 1063.19 460.68 191.48 229.18
48.76 1183.92 603.84 618.91 58.98 205.44 61.08 3085.36
AAREb p-CA 1.7 4.9 1.4 2.5 2.5 0.7 3.7 7.4
(Du12/R)/K
(Du21/R)/K
AARE
94.91 55.08 357.86 26.22 110.01 259.57 683.59 965.55
185.54 7.08 182.4 131.15 7.20 106.32 382.91 399.18
1.7 5.1 1.3 0.1 2.4 0.9 0.2 2.4
127.13 62.17 12.21 331.36 201.67 138.01 76.51 177.36
4.99 284.08 191.95 437.15 4.43 52.85 24.44 1895.92
3.8 2.7 0.5 0.9 22.0 13.3 1.6 1.0
CA
a b
Nonrandomness of the NRTL model was set equal to 0.2. nP parameter o NP AARE ¼ 1=NP ; NP is the number of points. i¼1 jxexp xcalc j=xexp 100
3.4 2.7 0.5 0.9 21.3 12.8 1.5 2.9
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FIGURE 5. Solubilities of p-coumaric acid in: (a) bmimBF4, bmimTFO, omimBF4 and t-pentanol; (b) bmimPF6, omimPF6, bmimTF2N and ethyl acetate. Solid lines correspond to UNIQUAC calculations.
By neglecting the terms that include DfusCp, equation (5) becomes:
Dfus H T ln ðc2 x2 Þ ¼ 1 Tm RT
ð6Þ
5.1. Correlation with the UNIQUAC & NRTL models Two classical local composition activity coefficient models were used for the correlation of the solubility data, the UNIQUAC [43] and NRTL [44] models. Temperature-independent binary interaction parameters per binary mixture were determined for both models, by minimizing the following objective function (OF):
OF ¼
8 2 91=2 P x2 c2;calc > < NP = 1 > i x c 2 2
> :
NP
> ;
;
ð7Þ
where NP is the number of experimental points, x2 and c2 are the experimental solubility and the corresponding activity coefficient
FIGURE 6. Solubilities of caffeic acid in: (a) bmimBF4, bmimTFO, omimBF4 and tpentanol; (b) bmimPF6, omimPF6, bmimTF2N and ethyl acetate. Solid lines correspond to UNIQUAC calculations.
values, and c2,calc is the activity coefficient calculated with the NRTL or UNIQUAC equation at x = x2,exp. The interaction parameters determined for the two models along with the obtained errors in solubility calculations are presented in table 5, whereas solubility results are graphically presented in figures 5 and 6. It is concluded that very good and similar correlation of the experimental data is obtained by both models, except from the solubilities of caffeic acid in bmimTF2N and bmimTFO that vary strongly with temperature. In these cases temperature-dependent interaction parameters are needed for accurate correlation. Overall, UNIQUAC yields slightly better results than NRTL. Moreover, the effect of the DfusCp, (equation (5)) was examined. Due to the lack of experimental data the value of DfusCp was calculated using predicted C Lp (406 J mol1 K1 for p-coumaric acid and 432 J mol1 K1 for caffeic acid) with the group contribution method of Ruzicka and Domalski [45,46] and predicted C Sp (357 J mol1 K1 for p-coumaric acid and 309 J mol1 K1 for caffeic acid) with the power-law method of Goodman et al. [47]. Unfortunately, worse results were obtained for NRTL and
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E.I. Alevizou, E.C. Voutsas / J. Chem. Thermodynamics 62 (2013) 69–78 TABLE 6 Prediction results with original and modified UNIFAC-Dortmund; mole fraction solubilities, x, of p-coumaric (p-CA) and caffeic acid (CA) in organic solvents. Solvent
T/K
p-CA
CA
x
ARE
a
x
ARE
Original UNIFAC
a
t-Pentanol
303.05 307.85 312.55 317.35
0.0562 0.0609 0.0658 0.0712
24.4 21.3 18.3 15.1
0.0886 0.0927 0.0969 0.1014
381.4 319.3 266.9 243.7
Ethyl acetate
303.05 307.85 312.55 317.35
0.0561 0.0618 0.0678 0.0744
279.1 276.8 289.7 289.5
0.0863 0.0915 0.0969 0.1027
5293.1 4981.7 4512.4 4365.2
t-Pentanol
303.05 307.85 312.55 317.35
Modified UNIFAC-Dortmund 0.0124 0.0158 0.0199 0.0249
83.3 79.6 75.3 70.3
0.0142 0.0182 0.0228 0.0283
22.9 17.8 13.7 4.2
Ethyl acetate
303.05 307.85 312.55 317.35
0.0005 0.0006 0.0008 0.0010
96.6 96.1 95.4 94.6
0.0694 0.0753 0.0813 0.0879
4240.0 4082.2 3773.3 3721.3
ARE ¼ jxexp xcalc j=xexp 100.
UNIQUAC when the DfusCp terms were taken into account that may be due to the predicted DfusCp values used.
5.2. Prediction of CADs solubility in Organic Solvents with UNIFAC The original UNIFAC model [48] and the modified UNIFAC-Dortmund [49] were applied for the prediction of CADs solubility only in tert-pentanol and ethyl-acetate, since UNIFAC parameters are missing for ILs. The results are presented in table 6. Original UNIFAC erroneously predicts higher solubilities of p-coumaric acid in ethyl acetate than in t-pentanol, except from the lowest temperature, while it correctly predicts higher solubilities of caffeic acid in t-pentanol than in ethyl acetate, except from the highest temperature. On the other hand, modified UNIFAC-Dortmund correctly predicts the ranking of the two solvents for the case of p-coumaric acid, while it predicts the opposite trend for caffeic acid. Nevertheless, both models give in most cases very poor qualitative predictions.
6. Conclusions The solubilities of p-coumaric acid and caffeic acid have been measured in six alkyl-substituted imidazolium-based ILs composed of the BF4, PF6, TF2N and TFO anions and in two organic solvents; t-pentanol and ethyl acetate. The results showed that the BF4 and TFO – based ILs are better solvents than the organic ones, which in turn are better than the corresponding PF6 and TF2N – based ILs. Furthermore, it is noticed that the increase of the alkyl chain length on the cation leads to a decrease in acids’ solubility in case of hydrophilic ILs with BF4 anion, but to an increase in case of hydrophobic ILs with PF6 anion. Moreover, melting temperatures and heats of fusion of the two acids were determined by differential scanning calorimetry, while thermogravimetric analysis revealed that both acids decompose forthwith melting. By utilizing the experimental solubilities, the thermodynamic functions of solutions were determined. The results indicated that the solution process is not spontaneous, is always endothermic and the entropy of solubilization is favorable for all systems studied. Finally, successful correlation with the NRTL and UNIQUAC models was obtained for all systems studied, with UNIQUAC being overall slightly
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JCT 12-431