Experimental determination and correlation of acetaminophen solubility in aqueous solutions of choline chloride based deep eutectic solvents at various temperatures

Fluid Phase Equilibria 462 (2018) 100e110

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Experimental determination and correlation of acetaminophen solubility in aqueous solutions of choline chloride based deep eutectic solvents at various temperatures Hemayat Shekaari*, Mohammed Taghi Zafarani-Moattar, Masumeh Mokhtarpour Department of Physical Chemistry, University of Tabriz, Tabriz, Iran

a r t i c l e i n f o

a b s t r a c t

Article history: Received 24 October 2017 Received in revised form 15 January 2018 Accepted 17 January 2018 Available online 31 January 2018

The aqueous solubility of acetaminophen (ACP) in some choline chloride (ChCl) based deep eutectic solvents (DESs) with urea, oxalic acid and malonic acid as neoteric green solvents were measured up to 0.90 wt fraction of DESs at T ¼ (293.15e318.15) K and at atmospheric pressure. The solubility in these solvents increased more than fortyefold with increasing the weight fraction of DESs, especially in DES containing malonic acid. The solubility data were accurately correlated by the NRTL, Wilson and UNIQUAC activity coefficient models. Also, to describe the thermodynamic behavior of ACP in the aqueous DES solutions, the thermodynamic functions, Gibbs energy, enthalpy, and entropy of dissolution and mixing were obtained by using the van't Hoff and Gibbs equations. The results indicate that the main contribution to the solubility of ACP in the aqueous DES solutions is enthalpic. © 2018 Elsevier B.V. All rights reserved.

Keywords: Deep eutectic solvent Solubility Acetaminophen Activity coefficient model Dissolution thermodynamic

1. Introduction The solubility in pharmaceutical industry is an paramount parameter to achieve desired concentration of drug in systemic circulation for pharmacological response to be shown. Drugs that have low solubility in water often show low bioavailability. Currently, only 8% of new drugs have both high solubility and permeability [1]. Various methods have been developed to enhance the solubility and bioavailability of the drugs including addition of surface-active agents, cyclodextrins, co-solvents, and pH adjustment. It is well-known that the addition of an organic co-solvent to water can significantly change the solubility of drugs [2]. Traditionally, organic solvents [3e6] and then after, in recent years ionic liquids (ILs) [7e9] have been used as co-solvent to improve the solubility of drugs but these types of solvents have problems including toxicity, flammability and high prices [10,11]. Therefore, developing a simple synthetic and greener alternative solvent has very significant practical importance. Recently, to overcome the limitations of organic solvents and ILs, deep eutectic solvents (DESs) have been developed in this area [12]. These neoteric solvents have advantages such as a lower cost and desirable

* Corresponding author. E-mail address: [email protected] (H. Shekaari). https://doi.org/10.1016/j.fluid.2018.01.017 0378-3812/© 2018 Elsevier B.V. All rights reserved.

environmental impact rather than the ILs and organic solvents [13]. In addition, they can be also prepared from biodegradable and natural components, and it was found that their toxicity is much lower than ILs [14]. These types of solvents are liquid at room temperatures typically formed by mixing two safe and relatively inexpensive solid compounds, such as a quaternary ammonium salt as hydrogen bond acceptor (HBA) (e.g. choline chloride (ChCl)) and a hydrogen bond donor (HBD) (e.g. urea or a carboxylic acid) at their eutectic composition with melting point much lower than that of the individual components [12]. The properties of DESs can be easily controlled by changing the mixing ratio of the HBA and HBD. Acetaminophen (ACP) is one of the most popular and most commonly used analgesic and antipyretic drugs around the world, available in mono- and multi-component preparations. The main issue with this drug is its low solubility in water (14 g L1 at 298.15 K), ACP is specially indicated in the treatment of several minor diseases presented by pediatric patients [15]. To the best of our knowledge, a few researches have been reported about the solubility of drugs in DES systems, which show increase in the solubility at higher concentrations of DESs [16e18]. Lu and coworkers [18] studied several drugs as models to explore the possibility of DESs application to improve their solubility and stability for potential nonaqueous liquid administration. The results indicate

H. Shekaari et al. / Fluid Phase Equilibria 462 (2018) 100e110

noticeable increasing in the solubility of ACP in pure DES including ChCl/urea, ChCl/oxalic acid and ChCl/malonic acid. In continuation of our previous work [19] the purpose of this work is providing the aqueous solubility of ACP in the presence of some DESs based on ChCl as HBA and urea, malonic acid, oxalic acid, as HBD up to 0.90 wt fraction of DESs at T ¼ (293.15e318.15) K and at atmospheric pressure. To correlate the solubility data, different activity coefficient models are suggested in literature. The local composition activity models such as Wilson [20], NRTL [21], and UNIQUAC [22] have been used in this study. Also, to describe the thermodynamic behavior of ACP in the aqueous DES solutions, thermodynamic functions, Gibbs energy, enthalpy, and entropy of dissolution and of mixing were obtained from the solubility data by using the van't Hoff and Gibbs equations at temperatures ranging from 293.15 K to 318.15 K.

101

features of these solvents are given in Table 2 and compared with those values in literature. The density, d, and speed of sound, u, of samples were measured with a vibrating tube densimeter (Anton Para, DSA 5000 densimeter and speed of sound analyzer) and the device was calibrated by distilled water. Also, refractive indices, nD, of the studied DESs were measured using a digital refractometer (ATAGO-DRA1, Japan). The apparatus was calibrated with doubly distilled water before each series of measurements. In addition, to ensure that the refractometer is working correctly, calibration was conducted with pure liquids of known refractive index such as hexane. The temperature was controlled using a circulating bath thermostat (Cooling Bath 490, Iran) with a thermal stability of ±0.1 K.

2.3. Solubility measurement 2. Experimental 2.1. Chemical The origin, CAS number and purity of the used chemicals are given in Table 1. The doubly distilled deionized water was used to prepare the solutions with a specific conductivity less than 1 mS cm1. 2.2. DES preparation The purified compounds of ChCl as HBA and urea, oxalic acid and malonic acid as HBDs were mixed with the molar ratio 1:2, 1:1, 1:1, respectively. The eutectic mixtures were stirred at 353.15 K for 4 h until a homogeneous, colorless liquid formed. Some of the

The binary solvent mixtures (DES þ water) were prepared by mixing the appropriate amounts (in grams) of the solvents with an analytical balance with precision 1  104 g (AW 220, GR220, Shimadzu, Japan). There are different methods to measure the solubility in the literature [26]. Among them, the shake flask method has been employed in this work. Briefly, the sealed vials containing an excess amount of ACP powder in the solvent mixtures were mixed using a shaker (Behdad, Tehran, Iran) and placed in water bath thermostat that was equipped with a temperaturecontrolling system with the uncertainty of 0.1 K (ED, Julabo Co., Germany) for 3 days to reach an equilibrium, the comparative calibration method was used to calibrate the thermostat, this method is carried out by comparing the thermostat with a higherquality reference thermometer. When a saturated solution was attained, the solid phase was removed by centrifugation using a

Table 1 Descriptions of the used chemicals. Chemical name

Provenance

CAS No.

Mass fraction (purity)

Acetaminophen

Merck

103-90-2

>0.98

Choline Chloride

Merck

67-48-1

>0.99

Urea

Merck

57-13-6

>0.98

Oxalic Acid

Merck

144-62-7

>0.99

Malonic Acid

Merck

141-82-2

>0.99

Structure

Table 2 Common properties of DESs used in this work at 298.15 K and 0.0866 MPa.a DES

ChCl/urea ChCl/malonic acid ChCl/oxalic acid a

Salt - HBD (molar ratio)

1:2 1:1 1:1

Melting Point (K)

285.15 [23] 283.15 [23] 307.15 [23]

MDES (g mol1)

259.74 243.68 229.65

103 d/(kg m3) Exp

Lit

1.193926 1.251470 1.210926

1.1979 [24] 1.2500 [16] 1.2200 [16]

Standard uncertainties for u (d) ¼ 0.006 kg m3, u (u) ¼ 0.50 m s1, u (nD) ¼ 0.0002, u (T) ¼ 0.1 K and u (P) ¼ 0.0001 MPa.

u(m s1)

2062.27 1962.69 1925.00

nD Exp

Lit

1.5041 1.4887 1.4809

1.5044 [25] 1.4871 [25] 1.4868 [25]

102

H. Shekaari et al. / Fluid Phase Equilibria 462 (2018) 100e110

Fig. 1. ACP calibration curve.

Hettich D-7200 centrifuge followed by filtration using Durapore® membrane filters, 0.45 mm. The clear solutions were diluted with water and assayed by a double beam UVevis spectrophotometer (Shimadzu, Japan) at 248 nm. The concentrations of the diluted solutions were determined from the calibration curve (Fig. 1) with the correlation coefficient of 0.9998. Each experimental data point represented the average of at least three repetitive experiments. The aqueous solubility of ACP in terms of drug mole fraction, x1 , in {ACP (1) þ water (2) þ DESs (3)}system was obtained by Eq. (1) [27]:

x1 ¼

w1 M1 w1 w2 M1 þ M2

(1)

w3 þM 3

where Mi and wi represent the molecular weight and mass fractions of i component in the saturated solution, respectively.

3.1. Modeling

0

11

0

(3)

j¼1

The knowledge of solubility is important in the pharmaceutical science, because it allows scientists and engineers to choose appropriate solvents for drug manufacturing processes. The solubility of a solid solute in a liquid solvent can be calculated by solving the thermodynamic equations of equilibrium. Modern theoretical development of equations that use the molar excess Gibbs energy, Gex , is often based on the concept of local composition theory, which is presumed to account for the short-range order and nonrandom molecular orientations that result from differences in molecular size and intermolecular forces. The solubility of a solute in a solution at a given temperature is calculated through the equality of the activity of the solute in the saturated solution and its activity in the pure solid state. Thus, through a solid-liquid equilibrium (SLE) framework, the following equation is obtained as [28]:

Dfus H 1 R

3.1.1. Wilson model The Wilson's equation gives not only an expression for the activity coefficients as a function of composition but also the variation of the activity coefficient with temperature. It provides a good representation of excess Gibbs energies, Gex , for partially miscible mixtures, such as solutions of polar and associating components (alcohols) in nonpolar solvents. The Wilson model for a solution with n-component was expressed the activity coefficient as [20]: n  n B CC BX  X B B Lki xk CC ln gi ¼ 1  lnB Lij xj  CC B n   P AA @ @ j¼1 k¼1 Lkj xj

3. Correlation procedure

ln x1 ¼

independent. To correlate the solubility data of the present drug in the aqueous DES solutions, one needs to take into account the nonideality of the liquid solution through an accurate activity coefficient function. Experimental activity coefficients were calculated by using Eq. (2) for the solutions. In order to find the proper model, all modeling were done for the nonelectrolyte and electrolyte systems. Once, the investigated systems were assumed as nonelectrolyte (ACP, water, DES) and the multicomponent Wilson, NRTL and UNIQUAC models were used to correlation of activity coefficient, and otherwise systems considered to be electrolyte (ACP, water, HBD, cholonium cation and chloride anion) also the Wilson, eeNRTL and UNIQUAC models for short range contribution and the PitzereDebyeeHückel (PDH) [29] model for long range contribution were applied. Since in appropriate modeling the number of parameters should be low and because in nonelectrolyte and electrolyte systems there are 6 and 12 parameters, respectively and the models average relative deviation percent (ARD%) were also lower in the non-electrolyte than the electrolyte models, we consider the system as non-electrolyte. In similar previous work, Gjineci and co-workers considered the system containing ethanol, water and DES as a nonelectrolyte system and DES in their work is a separate molecule [30]. In our calorimetric research lab, modeling calculations were performed with the mathcad software (PTC Mathcad® 15.0). The applied models in this system are as follows:

T



1 Tfus

where Lij is the binary interaction parameter which are related to the pure-component molar volumes, y, and to characteristic energy, l, differences by:



Lij ¼

yj lij  lii exp  yi RT

(2)

where Tfus , Dfus H, T, x1 and g1 refer to: melting temperature for the pure ACP, enthalpy of fusion for the pure ACP, (solid þ liquid) equilibrium temperature, equilibrium mole fraction, and the activity coefficient of the ACP in the saturated solution, respectively. The enthalpy of melting is assumed to be temperature

(4)

3.1.2. NRTL model Renon and Prausnitz [21] used the concept of local composition in the derivation of NRTL (Nonrandom Two-Liquid) equation. This equation is applicable to partially miscible as well as completely miscible systems. The activity coefficient in NRTL model is defined as follows [21]:

! þ ln g1



m P

ln gi ¼

tij Gij xj

j¼1 m P i¼1

Gli xl

0

1 x t G r rj rj C m X xj Gij B B C r¼1 þ Btij  m C m P P @ A j¼1 Glj xl Glj xl l¼1

m P

(5)

l¼1

where Gij ¼ expðaij tij Þ, aij ¼ aji and tii ¼ tjj ¼ 0. In this equation tij is the binary interaction parameter which is expressed by the following equation [21]:

H. Shekaari et al. / Fluid Phase Equilibria 462 (2018) 100e110

tij ¼

gij  gii RT

(6)

OF ¼

n  X

103

ln gexp  ln gcal i i

2

(13)

i¼1

where gij is an energy parameter characteristic of the iej interaction. The adjustable interaction parameter, Dgij, for this model is defined with Dgij ¼ gij  gii .

3.1.3. UNIQUAC model UNIQUAC (the Universal Quasi-Chemical theory) for ln gi consists of two parts: on the one hand, a combinatorial part, ln gCi , which describes the entropic contribution, and on the other hand a residual part, ln gRi , which accounts for the intermolecular forces that are responsible for the enthalpy of mixing. The combinatorial part depends on the composition, the size and the shape of the molecules. The residual part depends on intermolecular forces. The UNIQUAC equation [22] contains adjustable interaction parameters and is defined as:

ln gi ¼ ln gCi þ ln gRi

(8)

0

(9) 1

1

m m BX B C X qi tji B B qi tji C ln gRi ¼ qi B1 þ lnB A m P @ @ j¼1

qk tkj

j¼1

C C C A

qi xi ; m P q j xj

Fi ¼

j¼1



ri xi m P r j xj

(11)

j¼1

ln tij ¼

 

Duij

The temperature dependence of the solubility allows to insight into the molecular mechanisms involved in the solution processes using thermodynamic functions of dissolution [33]. In this study, the thermodynamic functions in the process of ACP dissolution are calculated based on the solubility of ACP in water and aqueous DES solutions as a function of temperature. The standard molar ο , is calculated from van't Hoff equaenthalpy of dissolution, DHsoln tion and defined as [33e35]: ο DHsoln ¼ R

The variables Fi , qi , and tji are the volume fraction, area fraction, and interaction parameter between molecule i and j, respectively. The coordination number, Z, the number of molecules surrounding the central molecule, is set to 10. Parameters r and q are pure component molecular-structure constants depending on molecular size and external surface areas. These parameters in this study are listed in Table 3. The adjustable interaction parameter which is related to an energy parameter characteristic of the iej interaction, Duij , for this model is:



where xexp , xcal and N are experimental and calculated solubility i i data and the number of experimental data, respectively.

(10)

k¼1

qi ¼

(14)

3.2. Thermodynamic properties of dissolution

z ðr  qi Þ  ðri  1Þ 2 i 0

1 N xexp xcal P ji i j B jxexp j C i C B ARD ¼ 100Bi¼1 C A @ N 0

(7)

    m F z q F X ln gCi ¼ ln i þ qi ln i þ li  i xl 2 xi Fi xi j¼1 j j lj ¼

exp

where n is the experimental points, also ln gi and ln gcal are i representing the experimental and calculated activity coefficients. The difference between the experimental and the calculated solubility data is defined by average relative deviation percent (ARD %) which is calculated for all activity coefficient models using the following equation:

 (12)

RT

The interaction parameters of the Wilson, NRTL and UNIQUAC models were determined by minimizing the objective function Eq. (13).

Table 3 UNIQUAC r and q parameters for the used components. Component

r

q

ACP [30] Water [31] ChCl/urea [32] ChCl/oxalic acid ChCl/malonic acid

5.4749 0.9200 9.9224 10.3836 11.0580

5.1801 1.4000 9.1200 9.0820 9.6220

vln x1   v 1=T

! (15) P

where x1 is the mole fraction of ACP solubility, R represents the universal gas constant (8.314 J K1 mol1) and T is the absolute temperature. The standard molar enthalpy change of solution, ο DHsoln , is generally obtained from the slope of the solubility curve in a so-called van't Hoff plot where lnx1 is plotted against T 1. Over a limited temperature interval, the heat capacity change of a solution may be assumed to be constant, hence the derived values of ο DHsoln will also be valid for the mean temperature, Tm ¼ 305.41 K and Eq. (15) can also be written as [36]:

1

0 B

ο DHsoln ¼ RB @ 

vln x1

v =T 1

1

=Tm

C C A

(16)

P

The standard molar Gibbs energy of the dissolution process, DGοsoln , can be calculated according to [37]:

DGοsoln ¼ RTm  intercept

(17)

where the intercept used is that obtained in plots of lnx1 versus (1/T - 1/Tm). The standard molar entropy of dissolution is also obtained from the following equation [33]:

DSοsoln ¼

ο DHsoln  DGοsoln

Tm

(18)

The xH and xTS represent the comparison of the relative contributions to the standard molar Gibbs energy by enthalpy and entropy in the dissolution process, respectively are expressed as follows [38]:

104

H. Shekaari et al. / Fluid Phase Equilibria 462 (2018) 100e110

Table 4 b c The experimental (xexp )a and calculated (xcal 1 ) solubility of ACP in the aqueous DES solutions at different temperatures (T) and weight fractions of DES (w3) from NRTL, Wilson 1 and UNIQUAC models. T/K

exp

103 x1

NRTL model 103 xcal 1

Wilson model 100

xexp xcal 1 1 xexp 1

103 xcal 1

100

xexp xcal 1 1 xexp 1

UNIQUAC model 103 xcal 1

100

xexp xcal 1 1 xexp 1

ACP (1) þ water (2) þ ChCl/urea (3) w3 ¼ 0.0000 293.15 298.15 303.15 308.15 313.15 318.15

1.589 ± 0.006 1.839 ± 0.002 2.239 ± 0.002 2.545 ± 0.005 3.009 ± 0.008 3.599 ± 0.004

1.593 1.857 2.250 2.547 3.064 3.675

0.21 0.96 0.48 0.08 1.82 2.13

1.592 1.832 2.237 2.560 3.028 3.615

0.20 0.40 0.12 0.62 0.61 0.45

1.594 1.844 2.237 2.550 3.006 3.014

0.26 0.21 0.12 0.22 0.13 0.15

w3 ¼ 0.2000 293.15 298.15 303.15 308.15 313.15 318.15

3.577 ± 0.004 3.983 ± 0.003 4.956 ± 0.003 5.577 ± 0.003 7.423 ± 0.007 8.364 ± 0.004

3.547 3.835 4.891 5.577 7.103 7.965

0.83 3.71 1.33 0.01 4.31 4.77

3.574 3.955 4.927 5.448 7.449 8.321

0.09 0.69 0.59 2.32 0.35 0.52

3.561 3.973 4.929 5.555 7.403 7.432

0.45 0.25 0.54 0.41 0.28 0.12

w3 ¼ 0.4000 293.15 298.15 303.15 308.15 313.15 318.15

6.177 ± 0.004 6.646 ± 0.002 8.878 ± 0.003 9.393 ± 0.003 15.380 ± 0.009 15.784 ± 0.006

6.231 6.991 8.981 9.387 15.421 15.745

0.87 5.20 1.16 0.07 0.27 0.25

6.180 6.648 8.846 9.396 15.422 15.778

0.04 0.03 0.36 0.03 0.27 0.04

6.219 6.672 8.908 9.494 15.387 15.396

0.67 0.40 0.35 1.07 0.04 0.10

w3 ¼ 0.6000 293.15 298.15 303.15 308.15 313.15 318.15

11.829 ± 0.004 12.515 ± 0.009 16.455 ± 0.004 18.055 ± 0.002 27.077 ± 0.004 28.048 ± 0.001

11.770 12.109 16.587 18.037 28.467 30.181

0.49 3.24 0.81 0.10 5.13 7.61

11.912 12.354 16.551 18.682 27.491 28.706

0.71 1.29 0.59 3.47 1.53 2.35

11.843 12.524 16.401 17.947 27.043 27.091

0.12 0.08 0.32 0.60 0.12 0.05

w3 ¼ 0.8000 293.15 298.15 303.15 308.15 313.15 318.15

26.243 ± 0.001 28.879 ± 0.002 36.521 ± 0.005 50.886 ± 0.007 60.198 ± 0.008 64.153 ± 0.005

26.218 29.119 35.763 50.855 56.737 59.266

0.09 0.83 2.08 0.06 5.75 7.62

26.220 28.808 36.291 49.626 60.142 63.502

0.09 0.24 0.63 2.48 0.09 1.01

26.244 28.811 36.490 51.080 60.105 60.204

0.01 0.23 0.08 0.38 0.15 0.01

w3 ¼ 0.9000 293.15 298.15 303.15 308.15 313.15 318.15

52.584 ± 0.006 59.421 ± 0.002 63.480 ± 0.005 72.780 ± 0.003 83.439 ± 0.006 97.302 ± 0.007

52.629 59.330 64.250 72.747 85.824 100.663

0.09 0.15 1.21 0.04 2.86 3.45

52.663 59.025 63.443 73.282 84.051 97.817

0.15 0.67 0.06 0.69 0.73 0.53

52.684 59.510 63.464 72.781 83.441 83.434

0.19 0.15 0.03 0.00 0.00 0.01

ACP (1) þ water (2) þ ChCl/oxalic acid (3) w3 ¼ 0.0000 293.15 298.15 303.15 308.15 313.15 318.15

1.589 ± 0.006 1.840 ± 0.002 2.239 ± 0.002 2.545 ± 0.005 3.009 ± 0.008 3.599 ± 0.004

1.589 1.841 2.262 2.534 3.012 3.268

0.04 0.08 1.02 0.44 0.09 1.02

1.604 1.848 2.249 2.555 3.002 3.619

0.96 0.43 0.42 0.40 0.25 0.57

1.588 1.848 2.235 2.552 3.024 3.248

0.05 0.44 0.18 0.27 0.48 0.43

w3 ¼ 0.2000 293.15 298.15 303.15 308.15 313.15 318.15

4.158 ± 0.006 4.199 ± 0.003 4.836 ± 0.005 5.905 ± 0.005 8.161 ± 0.005 9.083 ± 0.006

4.168 4.216 4.723 6.032 8.182 8.823

0.23 0.40 2.34 2.14 0.26 2.86

4.016 4.055 4.853 5.886 8.065 9.002

3.41 3.43 0.35 0.33 1.17 0.89

4.105 4.169 4.818 5.895 8.066 9.000

1.29 0.72 0.37 0.18 1.16 0.91

w3 ¼ 0.4000 293.15 298.15 303.15 308.15 313.15 318.15

6.777 ± 0.005 7.036 ± 0.005 9.778 ± 0.006 12.410 ± 0.006 16.154 ± 0.006 18.485 ± 0.008

6.744 6.954 9.865 11.864 15.824 18.715

0.48 1.16 0.90 4.40 2.04 1.24

6.820 7.028 9.831 12.293 15.942 18.461

0.64 0.10 0.54 0.94 1.31 0.13

6.934 7.195 9.783 12.453 16.427 18.787

2.32 2.27 0.05 0.35 1.69 1.63

H. Shekaari et al. / Fluid Phase Equilibria 462 (2018) 100e110

105

Table 4 (continued ) T/K

exp

103 x1

NRTL model 103 xcal 1

Wilson model 100

xexp xcal 1 1 xexp 1

103 xcal 1

100

xexp xcal 1 1 xexp 1

UNIQUAC model 103 xcal 1

100

xexp xcal 1 1 xexp 1

ACP (1) þ water (2) þ ChCl/urea (3) w3 ¼ 0.6000 293.15 298.15 303.15 308.15 313.15 318.15

14.162 ± 0.004 15.407 ± 0.005 20.056 ± 0.003 25.044 ± 0.007 25.720 ± 0.009 33.404 ± 0.009

14.234 15.561 20.284 26.344 26.932 34.821

0.51 1.00 1.14 5.19 4.71 4.24

14.955 16.023 20.119 25.594 26.438 34.381

5.60 4.00 0.31 2.20 2.79 2.92

13.863 15.172 20.032 25.087 25.659 33.181

2.11 1.52 0.12 0.17 0.24 0.67

w3 ¼ 0.8000 293.15 298.15 303.15 308.15 313.15 318.15

44.952 ± 0.005 46.210 ± 0.003 55.380 ± 0.008 56.521 ± 0.006 61.103 ± 0.006 75.012 ± 0.007

44.884 45.928 54.474 54.226 56.781 69.712

0.15 0.61 1.63 4.06 7.07 7.07

43.202 44.678 55.645 55.608 59.828 73.537

3.89 3.31 0.48 1.62 2.09 1.97

45.258 46.572 55.319 56.384 60.610 74.938

0.68 0.78 0.11 0.24 0.81 0.10

w3 ¼ 0.9000 293.15 298.15 303.15 308.15 313.15 318.15

62.988 ± 0.003 66.131 ± 0.005 79.029 ± 0.007 81.917 ± 0.007 86.145 ± 0.006 103.052 ± 0.008

63.058 66.250 79.758 83.311 89.799 107.234

0.11 0.18 0.92 1.70 4.24 4.06

63.739 66.721 79.316 81.979 86.539 103.727

1.19 0.89 0.36 0.08 0.46 0.66

62.893 66.079 79.070 82.045 86.663 103.261

0.15 0.08 0.05 0.16 0.60 0.20

ACP (1) þ water (2) þ ChCl/malonic acid (3) w3 ¼ 0.0000 293.15 298.15 303.15 308.15 313.15 318.15

1.589 ± 0.006 1.840 ± 0.002 2.239 ± 0.002 2.545 ± 0.005 3.009 ± 0.008 3.599 ± 0.004

1.591 1.887 2.224 2.531 3.022 3.077

0.09 2.58 0.69 0.52 0.42 0.37

1.604 1.845 2.249 2.539 3.004 3.614

0.97 0.28 0.44 0.23 0.19 0.43

1.591 1.837 2.230 2.534 3.011 3.058

0.10 0.15 0.41 0.41 0.07 0.25

w3 ¼ 0.2000 293.15 298.15 303.15 308.15 313.15 318.15

4.368 ± 0.006 4.844 ± 0.008 5.392 ± 0.008 6.223 ± 0.008 10.661 ± 0.001 12.268 ± 0.008

4.309 4.607 5.535 6.337 10.383 11.925

1.34 4.88 2.65 1.83 2.61 2.79

4.224 4.708 5.371 6.228 10.410 12.127

3.30 2.81 0.40 0.09 2.35 1.15

4.424 4.869 5.534 6.321 10.752 12.341

1.29 0.52 2.63 1.57 0.86 0.59

w3 ¼ 0.4000 293.15 298.15 303.15 308.15 313.15 318.15

8.539 ± 0.009 10.688 ± 0.008 13.116 ± 0.006 13.692 ± 0.009 16.691 ± 0.009 21.704 ± 0.007

8.601 10.553 12.467 13.365 17.340 22.422

0.72 1.27 4.95 2.39 3.89 3.31

8.514 10.612 12.981 13.593 16.624 21.649

0.30 0.71 1.03 0.73 0.40 0.26

8.101 10.248 12.352 13.306 16.177 21.133

5.13 4.13 5.82 2.82 3.08 2.63

w3 ¼ 0.6000 293.15 298.15 303.15 308.15 313.15 318.15

15.650 ± 0.008 18.264 ± 0.006 22.506 ± 0.005 25.195 ± 0.006 27.092 ± 0.009 39.779 ± 0.008

16.389 19.932 23.799 25.642 27.392 40.497

4.72 9.13 5.75 1.77 1.11 1.81

16.761 19.065 23.168 25.261 28.092 41.245

7.10 4.39 2.94 0.26 3.69 3.68

17.304 19.676 24.039 25.778 28.652 41.555

10.57 7.73 6.81 2.31 5.76 4.47

w3 ¼ 0.8000 293.15 298.15 303.15 308.15 313.15 318.15

51.842 ± 0.006 57.488 ± 0.007 58.866 ± 0.008 59.515 ± 0.008 68.239 ± 0.008 89.956 ± 0.006

44.745 52.389 56.283 59.208 61.548 83.020

13.69 8.87 4.39 0.51 9.80 7.71

49.487 55.465 58.090 59.360 66.275 87.917

4.54 3.52 1.32 0.26 2.88 2.27

45.679 52.241 56.145 58.759 64.249 85.649

11.89 9.13 4.62 1.27 5.85 4.79

w3 ¼ 0.9000 293.15 298.15 303.15 308.15 313.15 318.15

74.444 ± 0.007 79.680 ± 0.009 91.499 ± 0.004 99.788 ± 0.008 104.630 ± 0.008 140.519 ± 0.008

82.468 83.426 93.522 99.954 112.605 148.053

10.78 4.70 2.21 0.17 7.62 5.36

75.521 80.437 91.829 99.748 105.124 141.034

1.45 0.95 0.36 0.04 0.47 0.37

79.318 83.912 93.266 100.180 107.580 143.458

6.55 5.31 1.93 0.39 2.82 2.09

a b c

Standard uncertainty u(xexp 1 ) ¼ 0.5%. Standard uncertainty u(T) ¼ 0.1 K. Standard uncertainty u (w3) ¼ 0.0005.

106

H. Shekaari et al. / Fluid Phase Equilibria 462 (2018) 100e110

Fig. 2. The relationship between solubility of ACP, mole fraction x1, versus weight fraction of DES, wDES, and temperature in aqueous ChCl/urea solutions, the solid lines obtained from Wilson model.

 ο  DH   soln    þ T DSο   100 soln soln

(19)

  T DSο     100 ¼  ο  soln DHsoln þ T DSοsoln 

(20)

%xH ¼  ο DH %xTS

Fig. 4. The relationship between solubility of ACP, mole fraction x1, versus weight fraction of DES, wDES, and temperature in aqueous ChCl/malonic acid solutions, the solid lines obtained from Wilson model.

Tm ο ο DHsoln ¼ DHfus þ DHmix

(21)

m DSοsoln ¼ DSTfus þ DSοmix

(22)

where the thermodynamic functions of fusion process were calculated at mean temperatures as follows: 3.3. Thermodynamic functions of mixing T

The dissolution process may be represented by the following hypothetic steps [39], Solutesolid / Soluteliquid / Solutesolution To calculate the partial thermodynamic contributions to overall dissolution process, an approximation was used and then two equations were considered according to the literature [40].

T



fus m DSTfus ¼ DSfus  DCP ln

Tfus Tm



(23)

 (24)

where Table 5 The parameters of NRTLa activity coefficient model for the ACP in the aqueous DES solutions. T/K

Fig. 3. The relationship between solubility of ACP, mole fraction x1, versus weight fraction of DES, wDES, and temperature in aqueous ChCl/oxalic acid solutions, the solid lines obtained from Wilson model.



Tm fus DHfus ¼ DHfus  DCP Tfus  Tm

ACP (1) 293.15 298.15 303.15 308.15 313.15 318.15 ACP (1) 293.15 298.15 303.15 308.15 313.15 318.15 ACP (1) 293.15 298.15 303.15 308.15 313.15 318.15 a b

103Dgdwb

104Dgwd

104DgbdD

þ water (2) þ ChCl/urea (3) 3.94 1.36 4.42 4.03 1.39 0.68 3.83 1.35 6.62 7.13 2.64 0.80 3.51 1.31 0.98 3.33 1.28 1.04 þ water (2) þ ChCl/oxalic acid (3) 7.24 2.66 0.87 7.29 2.71 0.85 2.28 1.03 0.97 2.81 1.16 7.63 4.32 1.51 10.70 3.57 1.36 10.67 þ water (2) þ ChCl/malonic acid (3) 4.30 1.45 10.88 2.46 1.05 7.01 3.36 1.25 7.58 3.40 1.29 7.81 5.05 1.72 10.58 4.46 1.59 10.53

103DgDd

104DgwD

104DgDw

8.05 7.51 8.30 1.14 7.24 7.61

5.34 1.24 7.59 0.06 0.11 0.01

1.39 1.37 1.41 0.09 1.05 1.11

11.80 11.73 7.46 8.49 8.51 8.68

0.12 0.07 0.01 8.66 11.74 11.76

0.76 0.84 0.85 1.14 1.50 1.34

9.48 9.14 9.28 8.95 9.34 9.93

11.82 7.99 8.62 8.33 11.64 11.69

1.63 1.12 1.35 1.34 1.75 1.66

The non-randomness parameter of the NRTL model was set equal to 0.3. d ¼ drug, D ¼ DES, w ¼ water.

H. Shekaari et al. / Fluid Phase Equilibria 462 (2018) 100e110 Table 6 The parameters of Wilson model for the ACP in the aqueous DES solutions. 102Ldw

T/K

LdD

Lwd

ACP (1) þ water (2) þ ChCl/urea (3) 293.15 0.01 0.01 3.60 298.15 0.01 0.04 3.42 303.15 0.04 0.01 3.24 308.15 0.12 0.01 3.03 313.15 0.14 0.08 2.87 318.15 0.07 0.01 2.75 ACP (1) þ water (2) þ ChCl/oxalic acid (3) 293.15 0.09 0.03 3.53 298.15 0.11 0.01 3.33 303.15 0.16 0.07 3.13 308.15 0.12 0.03 3.02 313.15 0.06 0.01 2.91 318.15 0.01 0.03 2.72 ACP (1) þ water (2) þ ChCl/malonic acid (3) 293.15 0.08 0.005 3.53 298.15 0.14 0.04 3.29 303.15 0.10 0.02 3.21 308.15 0.03 0.01 3.08 313.15 0.03 0.01 2.92 318.15 0.07 0.003 2.76

LwD

LDd

LDw

0.10 0.70 0.34 0.01 0.11 0.16

4.54 41.54 8.16 3.42 2.14 3.23

3.76 13.27 2.90 1.37 0.84 1.53

0.06 0.06 0.23 0.02 0.01 0.13

4.55 4.36 5.86 3.10 4.08 2.46

1.56 1.51 1.75 1.00 1.72 0.88

0.004 0.01 0.16 1.12 0.01 0.36

6.57 3.39 5.81 33.48 16.20 22.79

2.11 0.81 2.34 6.01 6.16 9.81

103Dudw

ACP (1) 293.15 298.15 303.15 308.15 313.15 318.15 ACP (1) 293.15 298.15 303.15 308.15 313.15 318.15 ACP (1) 293.15 298.15 303.15 308.15 313.15 318.15

103DudD

103Duwd

103DuwD

þ water (2) þ ChCl/urea (3) 2.69 1.76 6.36 0.62 2.95 2.17 7.34 0.44 2.69 1.73 6.39 0.59 2.75 1.79 6.63 0.57 2.51 1.19 5.85 0.27 2.46 0.97 5.75 0.20 þ water (2) þ ChCl/oxalic acid (3) 2.85 0.91 6.91 0.43 2.78 0.77 6.68 0.38 2.88 0.82 7.05 0.28 2.64 0.30 6.24 0.37 2.62 0.11 6.20 0.27 2.65 0.06 6.34 0.19 þ water (2) þ ChCl/malonic acid (3) 2.46 2.03 5.63 0.84 2.29 1.55 5.16 2.61 1.34 25.63 3.11 0.30 0.47 1.94 1.87 0.84 2.60 1.95 6.12 0.72 2.37 1.95 5.51 0.71

103DuDd

103DuDw

3.26 4.75 3.37 3.68 2.47 2.21

0.19 0.15 0.21 0.22 0.49 0.58

4.30 4.21 4.60 3.63 3.27 3.60

0.15 0.06 0.07 0.06 0.07 0.09

0.82 1.96 0.94 1.01 0.90 1.01

3.35 4.32 2.11 1.11 3.45 3.26

T

DCP ¼

DHfusfus

(25)

Tfus

Table 8 The calculated average relative deviation percent (ARD%) for the solubility of the ACP in the aqueous DES solutions at several temperatures from different models. ARD%

Table 7 The parameters of UNIQUAC model for the ACP in the aqueous DES solutions. T/K

107

T

fus The values of DHfus and Tfus for ACP are 26.0 ± 0.2 kJ mol1 and 441.2 ± 0.5 K according to the literature [41]. The calculated values of enthalpy and entropy change of fusion at mean temperature from above mentioned equations are 17.99 kJ mol1 and 37.32 J mol1∙ K1, respectively.

T/K

NRTL

Wilson

UNIQUAC

0.02 0.55 0.39 1.60 0.59 0.08

0.28 0.22 0.24 0.44 0.12 0.07

0.54

0.23

2.62 2.03 0.41 0.92 1.34 1.18

1.00 0.96 0.15 0.23 0.83 0.65

Average 1.95 1.42 ACP (1) þ water (2) þ ChCl/malonic acid (3) 293.15 5.38 2.94 298.15 5.26 2.11 303.15 3.44 1.08 308.15 1.20 0.27 313.15 4.31 1.66 318.15 3.61 1.35

0.67

Average

3.54

ACP (1) þ water (2) þ ChCl/urea (3) 293.15 0.43 298.15 2.35 303.15 1.18 308.15 0.06 313.15 3.39 318.15 4.30 Average 1.95 ACP (1) þ water (2) þ ChCl/oxalic acid (3) 293.15 0.25 298.15 0.57 303.15 1.33 308.15 3.00 313.15 3.10 318.15 3.46

3.86

1.59

6.00 4.54 3.70 1.47 3.10 2.48

DESs has been shown in Figs. 2e4. It can be seen from these figures, the solubility of ACP was increased in the aqueous DES solutions to more than 40-fold at higher concentration of DES and temperatures. It seems that in the solutions, ACP acts as HBA and urea, malonic and oxalic acid, act as HBD. The eCOOH groups of oxalic acid interact strongly with ACP while is weak in the case of malonic acid. This turns possible a stronger interaction of malonic acid with ACP which leads to higher solubility of ACP in malonic acid based DES compared to oxalic acid and urea based DES. The results indicate that these types of neoteric green solvents are proper solvents rather than ILs and organic solvents in pharmaceutical fields. On the other hand, there are some reports on solubility of ACP in co-solvent systems. In methanol þ water co-solvent with weight fraction of 0.2 for methanol at 298.15 K, the value of ~ oz et al. [42]. 4.38  103 (mole fraction) has been reported by Mun The value of 4.84  10 3 obtained at the same temperature and weight fraction for solubility of ACP in DES co-solvent system containing malonic acid indicate that there is an improvement in nez et al. the solubility of this drug using this DES. According to Jime [40] solubility of ACP in propylene glycol þ water co-solvent mixtures at 298.15 K and weight fraction of 0.2 for propylene glycol, is 3.47  103, this value is lower than solubility we found in DES containing malonic acid. There are some cosolvent systems like ethanol þ water and dioxan in which slightly higher solubility of ACP have been reported [43,44]. 4.2. Modeling results

4. Results and discussion 4.1. Solubility results The experimental solubility data of ACP in binary solvents (DES þ water) with different weight fractions of DES at various temperatures (293.15e318.15 K) are listed in Table 4. The relationship between solubility of ACP, x1 , versus absolute temperature in the aqueous DES solutions with different weight fractions of

In the following, the solubility data of ACP in the aqueous solutions were correlated with the Wilson, NRTL and UNIQUAC models. The modeling results and the corresponding parameters are summarized in Tables 4e7. The calculated ARD% values are given in Table 8 for the models in this work. Thus, the performance of these models in correlation of the experimental solubility data can be ordered as UNIQUAC > Wilson > NRTL for systems containing urea an oxalic acid, also Wilson > UNIQUAC > NRTL for malonic acid

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H. Shekaari et al. / Fluid Phase Equilibria 462 (2018) 100e110

Table 9 The calculated activity coefficient of ACP, g1 , based on Wilson model in aqueous DES solutions at different temperatures. DES weight fraction

T ¼ 293.15 K

ACP (1) þ water (2) þ ChCl/urea (3) 0.0000 17.46 0.2000 7.78 0.4000 4.50 0.6000 2.33 0.8000 1.06 0.9000 0.53 ACP (1) þ water (2) þ ChCl/oxalic acid (3) 0.0000 17.33 0.2000 6.92 0.4000 4.08 0.6000 1.86 0.8000 0.64 0.9000 0.44 ACP (1) þ water (2) þ ChCl/malonic acid (3) 0.0000 17.33 0.2000 6.58 0.4000 3.27 0.6000 1.66 0.8000 0.56 0.9000 0.37

T ¼ 298.15 K

T ¼ 303.15 K

T ¼ 308.15 K

T ¼ 313.15 K

T ¼ 318.15 K

18.15 8.41 5.00 2.69 1.15 0.56

17.68 8.03 4.47 2.39 1.09 0.62

18.26 8.58 4.98 2.50 0.94 0.64

18.16 7.38 3.57 2.00 0.91 0.65

17.80 7.73 4.08 2.24 1.01 0.66

18.00 8.20 4.73 2.08 0.74 0.50

17.59 8.15 4.02 1.97 0.71 0.50

18.30 7.94 3.80 1.83 0.84 0.57

18.32 6.82 3.45 2.08 0.92 0.64

17.78 7.15 3.49 1.87 0.87 0.62

18.08 7.06 3.13 1.74 0.60 0.41

17.58 7.36 3.05 1.71 0.68 0.43

18.42 7.51 3.44 1.85 0.79 0.47

18.31 5.28 3.31 1.96 0.83 0.52

17.80 5.31 2.97 1.56 0.73 0.46

systems. The calculated activity coefficient values from the Wilson model are given in Table 9. From these values, a rough estimate of soluteesolvent intermolecular interactions can be made and the results show that g1 values decrease with increasing the weight fraction of DESs. The comparison of calculated solubility and activity coefficient from this model is shown in Fig. 5 which indicates that the drug desire to precipitate with increasing the activity coefficient of the drug in the solution, thus ACP-DES interactions will decrease. 4.3. Thermodynamic properties of dissolution and mixing results

Fig. 5. The relationship between calculated solubility and activity coefficient for ACP in the aqueous malonic acid DES solution from Wilson model at various temperatures; 293.15(-), 303.15 (A), 313.15 (:), 318.15 (C).

To calculate the Thermodynamic properties of dissolution, the   plots of ln x1 versus 1=T  1=T for ACP in aqueous DES solutions m

containing malonic acid were shown in Fig. 6. ο , and T DSο The results of DGοsoln , DHsoln are collected in m soln

Table 10 Thermodynamic functions for dissolution process at different weight fractions of DES (w3) w3

ο DHso /kJ mol1 ln

ACP (1) þ water (2) þ ChCl/urea (3) 0.0000 25.33 ± 0.15 0.2000 27.58 ± 0.08 0.4000 32.14 ± 0.28 0.6000 29.74 ± 0.16 0.8000 31.08 ± 0.14 0.9000 18.71 ± 0.18 ACP (1) þ water (2) þ ChCl/oxalic acid (3) 0.0000 25.20 ± 0.20 0.2000 26.91 ± 0.16 0.4000 34.41 ± 0.23 0.6000 26.58 ± 0.15 0.8000 15.11 ± 0.24 0.9000 14.56 ± 0.19 ACP (1) þ water (2) þ ChCl/malonic acid (3) 0.0000 18.75 ± 0.32 0.2000 33.84 ± 0.11 0.4000 26.76 ± 0.21 0.6000 26.36 ± 0.21 0.8000 14.22 ± 0.17 0.9000 18.00 ± 0.11 a,b

Standard uncertainty u is u(w3) ¼ 0.0005, u(T) ¼ 0.1 K.

a

at mean temperature b.

Tm DSοso ln /kJ mol1

DGοso ln /kJ mol1

xH

xTS

9.99 ± 0.22 14.31 ± 0.13 20.36 ± 0.17 19.53 ± 0.12 23.03 ± 0.20 11.95 ± 0.15

15.35 ± 0.17 13.27 ± 0.22 11.78 ± 0.12 10.21 ± 0.22 8.05 ± 0.14 6.76 ± 0.13

71.73 65.83 61.22 60.36 57.44 61.02

28.27 34.17 38.78 39.64 42.56 38.98

9.85 ± 0.10 13.80 ± 0.12 22.94 ± 0.13 16.79 ± 0.12 7.77 ± 0.27 8.10 ± 0.20

15.35 ± 0.16 13.09 ± 0.15 11.46 ± 0.16 9.78 ± 0.10 7.34 ± 0.16 6.45 ± 0.12

71.90 66.09 60.00 61.28 66.04 64.25

28.10 33.91 39.99 38.71 33.95 35.75

3.39 ± 0.08 21.14 ± 0.09 15.82 ± 0.10 16.84 ± 0.12 7.208 ± 0.20 12.06 ± 0.18

15.36 ± 0.24 12.70 ± 0.13 10.94 ± 0.11 9.51 ± 0.13 7.01 ± 0.19 5.95 ± 0.14

84.68 61.55 62.85 61.01 66.35 59.89

15.32 38.45 37.15 38.99 33.65 40.11

H. Shekaari et al. / Fluid Phase Equilibria 462 (2018) 100e110

Fig. 6. Plot of lnx1 vs (1/T- 1/Tm); in aqueous ChCl/malonic acid solutions at different weight fraction of DES (wDES): 0.0000(A), 0.2000 (-), 0.4000 (:), 0.6000 (C), 0.8000 (,), 0.9000 (o).

Table 10. The standard molar Gibbs energy and enthalpy of dissolution are positive in all systems, therefore the process of ACP dissolution in the aqueous DES solutions is always endothermic. The DGοsoln values decrease with increasing the weight fraction of DES, which indicating that the solubility of ACP in these types of the solvents increases with the decrease of the DGοsoln values. On the other hand, DSοsoln is positive in all studied solutions and less than ο , which show that the enthalpic is as driving overall the DHsoln dissolution process for all the solutions. This is attributed to the possible breaking of highly ordered structure of solvent molecules surrounding the drug molecules. The calculated xH and xTS values are given in Table 10. From this table it follows that the main contribution to standard molar Gibbs energy of dissolution process of ACP is the enthalpic (greater than 54% in all cases). The thermodynamic functions of mixing for solubility of ACP in the aqueous DES solutions were summarized in Table 11. By analyzing the partial contributions by ideal solution (related to solute fusion process) and mixing processes to the enthalpy and 305:5 and DS305:5 are positive. entropy of solution, it is found that DHfus fus

Table 11 Thermodynamic functions for mixing process at different weight fractions of DES (w3) a at mean temperature b. w3

ο /kJ mol1 DHmix

Tm DSοmix /kJ mol1

ACP (1) þ water (2) þ ChCl/urea (3) 0.0000 7.34 ± 0.10 1.42 ± 0.09 0.2000 9.59 ± 0.17 2.91 ± 0.09 0.4000 14.14 ± 0.12 8.96 ± 0.09 0.6000 11.74 ± 0.10 8.12 ± 0.12 0.8000 13.08 ± 0.20 11.62 ± 0.12 0.9000 0.71 ± 0.17 0.54 ± 0.10 ACP (1) þ water (2) þ ChCl/oxalic acid (3) 0.0000 7.20 ± 0.16 1.55 ± 0.11 0.2000 8.91 ± 0.12 2.41 ± 0.12 0.4000 16.41 ± 0.15 11.54 ± 0.23 0.6000 8.58 ± 0.05 5.39 ± 0.16 0.8000 2.89 ± 0.12 3.63 ± 0.11 0.9000 3.43 ± 0.12 3.29 ± 0.27 ACP (1) þ water (2) þ ChCl/malonic acid (3) 0.0000 0.75 ± 0.15 8.01 ± 0.20 0.2000 15.84 ± 0.17 9.74 ± 0.15 0.4000 8.76 ± 0.13 4.41 ± 0.17 0.6000 8.36 ± 1.64 5.44 ± 0.12 0.8000 3.78 ± 0.19 4.19 ± 0.15 0.9000 0.01 ± 0.05 0.65 ± 0.11 a,b

DGοmix /kJ mol1

xH

xTS

8.76 ± 0.17 6.67 ± 0.15 5.18 ± 0.16 3.62 ± 0.18 1.46 ± 0.17 0.16 ± 0.05

83.81 76.70 61.21 59.10 52.95 87.19

16.18 23.29 38.78 40.89 47.04 12.81

8.75 ± 0.15 6.50 ± 0.25 4.87 ± 0.22 3.19 ± 0.16 0.74 ± 0.12 0.14 ± 0.08

82.32 78.72 58.72 61.41 44.32 51.05

17.67 21.27 41.28 38.59 55.67 48.95

8.76 ± 0.20 6.11 ± 0.11 4.35 ± 0.12 2.92 ± 0.13 0.41 ± 0.13 0.65 ± 0.13

8.61 61.93 66.50 60.57 47.41 0.67

91.38 38.06 33.50 39.43 52.59 99.33

Standard uncertainty u is u(w3) ¼ 0.0005, u(T) ¼ 0.1 K.

109

Fig. 7. The DGοmix values relative to mixing process of ACP in the aqueous DES solutions at 305.41 K, ChCl/urea (-), ChCl/oxalic acid (A), ChCl/malonic acid (:).

ο DHmix is positive in those mixtures containing DES and by increasing the amount of DES, it becomes negative. The net variaο tion in DHmix values results from the contribution of several kinds of interactions. The enthalpy of cavity formation is endothermic because energy must be supplied to overcome the cohesive forces of the solvent this process decreases solubility [40]. On the other hand, the enthalpy of solute-solvent interaction is exothermic and it is originated mainly from the van der Waals and Lewis acid-base interactions [40]. The structuring of water molecules around the nonpolar groups of solutes (hydrophobic hydration) contributes to decrease the net heat of mixing to small or even negative values in aqueous solutions. According to Romero et al. [5] in the initial portion of the solubility curve, the hydrogen bonding of ACP will increases with co-solvent concentration. The entropy of mixing, DSοmix , in water has a negative value and it is positive at low weight fraction of DES and it becomes small or negative at high weight fractions of DESs. DGοmix for studied systems has been shown in Fig. 7, according to that, the DGοmix of systems become small and negative with increasing the amount of DES, and it is more negative for the system containing malonic acid. In addition, the negative amount of DGοmix is a beneficial factor for dissolving, and solubility of ACP in DES with malonic acid is higher than other investigated systems in this article.

5. Conclusions The aqueous solubility of acetaminophen in the presence of some deep eutectic solvents (ChCl/urea, ChCl/oxalic acid and ChCl/ malonic acid), as co-solvents was determined experimentally within the temperatures ranging from (293.15e318.15) K. The acetaminophen mole fraction of solubility in the studied solvents increased approximately 40-fold with increasing deep eutectic solvents concentration and temperature. Malonic acid based deep eutectic solvent exhibit maximum and solvent containing urea has minimum solubility. Moreover, the local composition activity coefficient models for nonelectrolyte systems such as Wilson, NRTL, and UNIQUAC were applied to correlate the solubility data. The results indicate that the performance of these models in correlation of the experimental solubility data can be ordered as UNIQUAC > Wilson > NRTL for systems containing urea an oxalic acid, also Wilson > UNIQUAC > NRTL for malonic acid systems. Also, thermodynamic functions of dissolution and mixing were calcuο lated for these systems and the positive DGοsoln and DHsoln suggest that the dissolution process of acetaminophen to be endothermic and nonspontaneous which this trend decreases with addition of ο DES in aqueous media. On the other hand DGοmix and DHmix values are positive but with addition of DES they become small or

110

H. Shekaari et al. / Fluid Phase Equilibria 462 (2018) 100e110

negative, and this result show that solubility of acetaminophen increases with decreasing of DGοmix . Also, the main contribution to standard molar Gibbs energy in dissolution and mixing process of ACP in the aqueous DES solutions is enthalpic.

[20]

[21]

Acknowledgements [22]

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