Article Cite This: J. Chem. Eng. Data XXXX, XXX, XXX-XXX
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Thermodynamic Aspects of Solubility and Solvation of Bioactive Bicyclic Derivatives in Organic Solvents Svetlana V. Blokhina,† Angelica V. Sharapova,*,† Marina V. Ol’khovich,† Tatyana V. Volkova,† Alexey N. Proshin,‡ and German L. Perlovich† †
Institute of Solution Chemistry, Russian Academy of Sciences, 1 Akademicheskaya Street, 153045, Ivanovo, Russia Institute of Physiologically Active Compounds, Russian Academy of Sciences, 142432, Chernogolovka, Russia
‡
S Supporting Information *
ABSTRACT: The solubility of eight biologically active nonaromatic compounds with a common 3-thia-1-aza-bicyclo fragment in 1-octanol and hexane was determined in the temperature range of T = 293.15−315.15 K. The solubility of the substances in alcohol is approximately 2 orders of magnitude higher than in alkane and varies within 3.0 × 10−3 to 2.1 × 10−1 and 7.1 × 10−5to 13.5 × 10−3 mol. fractions, respectively. In accordance with the structure of the substituent at the para position of the phenyl ring, the studied compounds are arranged in the following order of decreasing solubility in both solvents: methyl- > fluoro- > ethyl- > trifluoromethyl- > cyano- > acetyl-. The ideal solubility of the substances in 1-octanol, calculated from thermophysical parameters was found to correlate with the reported experimental data. The activity coefficients of the substances in saturated solutions of organic solvents are determined by the method of distribution between two liquid phases. The thermodynamic aspects of the relationship between the processes of dissolution, melting, sublimation, and solvation of the substances are discussed.
1. INTRODUCTION Heterocyclic compounds are widespread in nature and are the key components in biological processes.1 Therefore, a large number of works are devoted to the methods of synthesizing bioactive compounds of this class and to studying their pharmaceutically significant properties.2 It was found earlier that 1,3-thiazine derivatives have antimicrobial activity and can be used to treat neurodegenerative disorders, disorders of gastrointestinal motility, and inflammation.3,4 New representatives of this series are compounds based on bicycle[3.3.1] nonane core that possess radioprotective ability.5 The introduction of an additional cycle in the parent structure of thiazine promoted lipophilicity growth and, as a result, increased the biological activity of the obtained bridged dihydrothiazine bicycles, which are promising candidates for drugs. One of the main drug compound properties which determines their bioavailability, optimal therapeutic doses, and possible side effects is solubility in biological media.6 The solubility of most of the synthesized test substances does not satisfy the requirements of the Biopharmaceutics Classification System, which creates difficulties for their application.7 It is possible to solve the problem if the molecular fragments which improve the solubility and other important transport properties are revealed before the synthesis stage. The study of the regularities of the bioactive compounds covalent modification effect on their physicochemical properties is aimed at © XXXX American Chemical Society
optimizing the search for substances with improved pharmaceutical characteristics at the in vitro stage.8 The most important solution properties directly related to the intermolecular solute−solute and solute−solvent interaction are the thermodynamic parameters of dissolution and solvation. Nevertheless, the number of systematic studies of these processes for polycyclic compounds as a specific type of organic substances is limited in the modern literature. This predetermined the urgency of the task set in the present paper to establish general regularities and features of dissolution and solvation of biologically active derivatives of bicycle [3.3.1] nonane in nonaqueous media, depending on the chemical nature of the functional substituents and solvents. It is quite natural that at the first stage, our attention was focused on the aqueous solution, as water plays an important role in biological processes.9 In addition, we conducted earlier the sublimation experiments and determined the energies of the crystal lattices of the investigated compounds, which was necessary for constructing a thermodynamic cycle and calculating the solvation functions.10 1-Octanol and hexane were chosen as solvents in this work. The amphipathic nature of 1-octanol makes it possible to use it as a model solvent, best imitating the lipid layer of membranes and proteins in animal and plant Received: July 14, 2017 Accepted: October 20, 2017
A
DOI: 10.1021/acs.jced.7b00641 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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Table 1. Source, Purification, and Analysis Details of the Chemicals Used chemical name (4-methyl-phenyl)-[3-thia-1-aza-bicyclo[3.3.1]non-2ylidene]-amine (1) (4-ethyl-phenyl)-[3-thia-1-aza-bicyclo[3.3.1]non-2ylidene]-amine (2) (4-isopropyl-phenyl)-[3-thia-1-aza-bicyclo[3.3.1]non-2ylidene]-amine (3) (4-fluoro-phenyl)-[3-thia-1-aza-bicyclo[3.3.1]non-2ylidene]-amine (4) (4-trifluoromethyl-phenyl)-[3-thia-1-aza-bicyclo[3.3.1]non2-ylidene]-amine (5) (4-cyano-phenyl)-[3-thia-1-aza-bicyclo[3.3.1]non-2ylidene]-amine (6) (4-acetyl-phenyl)-[3-thia-1-aza-bicyclo[3.3.1]non-2ylidene]-amine (7) (3-chloro-4-methyl-phenyl)-[3-thia-1-aza-bicyclo[3.3.1] non-2-ylidene]-amine (8) 1-octanol hexane potassium dihydrogen phosphate disodium hydrogen phosphate dodecahydrate a
method purification
final mole fraction purity
1639369-07-5
recrystallization
≥0.95
1
synthesis
1583299-20-0
recrystallization
≥0.95
1
synthesis
1778735-15-1
recrystallization
≥0.95
1
synthesis
1639369-06-4
recrystallization
≥0.95
1
synthesis
1639369-05-3
recrystallization
≥0.95
1
synthesis
1778735-27-5
recrystallization
≥0.95
1
synthesis
1778736-36-9
recrystallization
≥0.95
1
synthesis
1583299-21-1
recrystallization
≥0.95
1
SigmaAldrich SigmaAldrich Merck Merck
111-87-5
≥0.99a
none
110-54-3
≥0.97a
none
7778-77-0 10039-32-4
≥0.99 ≥0.99a
none none
source
CAS number
synthesis
initial mole fraction purity
a
analysis method H NMR H NMR H NMR H NMR H NMR H NMR H NMR H NMR
As stated by the supplier.
cells.11 Hexane is an aprotic solvent capable only of van der Waals interactions, which allows us to obtain information on specific interactions of the solutes with the solvent based on comparison.
8) in the both solvents have been specified as 259, 280, 250, 260, 270, 298, 296, and 282 nm, respectively. The experimental results are reported as an average value of at least three replicated experiments. It should be noted that the sediment DSC analysis showed that there were no crystallosolvates in all the tested compounds. The standard solution enthalpies ΔH0solv were calculated using the van’t Hoff equation
2. EXPERIMENTAL SECTION Materials. The method of synthesis, 1H NMR spectroscopy, and elemental analysis data for compound studied have been reported earlier.9 Graphical 1H NMR spectra of compounds studied have been provided in Figure S1. Bidistilled water (with electrical conductivity 2.1 μS cm−1) was used to prepare the buffer solution. The phosphate buffer pH 7.4 (I = 0.15 mol· L−1) was prepared by combining the KH2PO4 (9.1 g in 1 L) and Na2HPO4·12H2O (23.6 g in 1 L) salts. The pH values were measured by using a pH meter FG2-Kit (Mettler Toledo, Switzerland) standardized with pH 1.68, 6.86, and 9.22 solutions. The detailed information about all the chemicals used in this work is listed in Table 1. Solubility. All the experiments were carried out by the isothermal saturation method at five temperature points: 293.15, 298.15, 303.15, 310.15, and 315.15 ± 0.1 K. The essence of the above method includes determination of the compound concentration in the saturated solution. Glass ampules containing the tested substance and the solvent were placed into the air thermostat equipped with a stirring device. The time required for reaching a constant value of the solution concentration was determined from the solubility kinetic dependences and averaged 24 h. The time of solid phase sedimentation after stirring was 2 h. The solution aliquot was taken and centrifugated in a centrifuge with the temperature controlled by Biofuge stratos (Germany) during 5 min under a fixed temperature. The solid phase was removed by isothermal filtration with the filter MILLEXHA 0.45 μm (Ireland). The saturated solution was diluted with the correspondent solvent to the required concentration. The molar solubilities of the drugs were measured by the spectrophotometer Cary-50 (Varian, U.S.A.) with an accuracy of 2−4%. The wavelengths corresponding to the absorption maximums of compounds (1−
0 ΔHsol d(ln x) = dt RT 2
(1)
The temperature dependences of drug solubilities within the chosen temperature interval can be described by the linear function ln x = A −
B T
(2)
This indicates that the change in heat capacity of the solutions with the temperature is negligibly small. Distribution. The distribution experiments were carried out according to the shake flask method12 in the following manner: 4 mL of octanol-saturated buffer pH 7.4 was put in a glass flask and carefully overlaid with 4 mL of buffer-saturated octanol which was already spiked with a definite amount of the test chemical. The glass-jacketed stirring flask was held at a constant temperature of 298.15 K by means of an air thermostat. The distribution equilibrium of the chemical between both phases was reached only after 24 h. In the actual investigation we took samples from both phases after reaching an equilibration. The compound concentration in both phases was determined spectrophotometrically according to the procedure described in Solubility. The reproducibility of the measured concentrations was below 0.1%, and the maximum deviations from the average value were always −F(4) > −CH2−CH3(2) > −CF3(5) > −CN(6) > −COCH3(7) > −CH3,-Cl(8). Compound (3) with an isopropyl substituent is an exception from the general tendency of the solubility variation. Compound (1) with a methyl substituent in the phenyl ring has the highest solubility among the other substances. The solubility of compound with ethyl substituent (2) is lower than that of compound with isopropyl substituent (3), which is due to the lower melting point of the latter. Interestingly, the differences in the structure of the alkyl substituent of compounds (1) and (3) have practically no effect on the solubility in 1-octanol. In contrast, the following regularity of the solubility in hexane for the compounds with alkyl substituents was revealed: ethyl < isopropyl < methyl. The substances (5−8) with halogens, nitrogen, and oxygen atoms belong to the group of compounds with low solubility. The experiment showed that the introduction of a chlorine atom into the meta- position of the phenyl ring of compound (1) leads to a decrease in the solubility of compound (8) in both solvents by almost 2 orders of magnitude. It should be noted that according to the data obtained earlier9 the solubility of compounds studied in a buffer solution pH 7.4 sighificantly
Figure 1. Molecular structure of the studied compounds.
solubility values of the substances studied in 1-octanol and hexane obtained by the isothermal saturation method are summarized in Tables 3 and 4. As it follows from the data presented, the solubility of the substances in both solvents increases with the temperature growth and varies within 3.0 × 10−3 to 2.1 × 10−1 and 7.1 × 10−5 to 13.5 × 10−3 mole fraction in the temperature range T = 293.15 to 315.15 K for solutions of 1-octanol and hexane, respectively. For all the studied bicycles the solubility in 1-octanol is higher than in hexane, which is explained by the differences in the intermolecular interaction of the solutes and solvents (Figure 2). The compounds studied herein have a tertiary nitrogen atom in the heterocycle structure, which is an electron donor center and determines their behavior in solutions as typical bases. The higher solubility of the substances in 1-octanol is due to the formation of a hydrogen bond between the solvent molecule, which acts as a proton donor, and the molecule of the dissolved proton acceptor substance. In this case, the
Table 2. Structural Formula of the Radical (R), Molecular Mass (M), Melting (Tm, ΔHm) and Sublimation (ΔG0sub ΔH0sub, ΔS0sub) Parameters 1of the Compounds Studied
a
The melting parameters have been determined by the authors earlier.9 bThe experimental values of the sublimation enthalpy have been determined by the authors earlier.13,14 C
DOI: 10.1021/acs.jced.7b00641 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
2.45 2.96 3.54 4.69 5.67 8.4 ± 0.2 3543 ± 71 0.998 2.41 × 10−2
2
1.18 1.32 1.46 1.71 1.89 2.4 ± 0.2 2013 ± 69 0.991 2.33 × 10−2
x × 10
5
The standard uncertainties are u(T) = 0.15 K, u(p) = 0.003 MPa. bThe relative standard uncertainty is ur(x) = 0.04.
14.3 15.4 16.4 18.2 19.4 2.43 ± 0.07 1282 ± 22 0.9995 0.4 × 10−2
x × 10
x × 10 2
4 2
3 2
0.87 1.07 1.31 1.66 2.02 7.19 ± 0.2 3496 ± 61 0.9995 1.1 × 10−2
x × 10
6
D
a
3.30 4.65 6.54 10.40 13.45 14.7 ± 0.4 5970 ± 107 0.9995 1.3 × 10−1
293.15 298.15 303.15 310.15 315.15 A B R σ
2.25 2.89 3.82 5.37 6.73 7.43 ± 0.10 4639 ± 32 0.9998 0.8 × 10−2
1.35 1.58 1.95 2.46 2.88 4.4 ± 0.24 3233 ± 74 0.9992 1.4 × 10−2
x × 10
x × 10 3
3 4
2 3
1.50 1.87 2.24 2.94 3.59 5.9 ± 0.2 3637 ± 54 0.9997 3.3 × 10−2
x × 10
4
4
2.48 2.86 3.39 4.02 4.60 0.6 ± 0.2 2597 ± 63 0.9991 4.32 × 10−2
x × 10
5
The standard uncertainties are u(T) = 0.15 K, u(p) = 0.003 MPa. bThe relative standard uncertainty is ur(x) = 0.04.
x × 10
3
T/K
1
4
2.23 2.71 3.37 4.22 5.23 3.64 ± 0.32 3531 ± 96 0.9989 1.8 × 10−2
x × 10
6
Table 4. Experimental Mole Fraction Solubility (x) of Compounds Studied in Hexane at Different Temperatures (p = 0.1 MPa)a,b
a
1.45 1.57 1.71 1.93 2.09 3.3 ± 0.1 1542 ± 19 0.999 3.85 × 10−2
293.15 298.15 303.15 310.15 315.15 A B R σ
12.4 13.9 16.3 19.4 22.4 4.07 ± 0.21 2482 ± 66 0.9989 1.2 × 10−2
x × 10
x × 10
T/K
3
2
1 2
x × 10
7
4
1.44 1.76 2.15 2.80 3.44 3.55 ± 0.16 3633 ± 50 0.9997 0.9 × 10−2
x × 10
7
0.75 0.93 1.18 1.56 2.02 9.07 ± 0.36 4094 ± 109 0.9989 2.1 × 10−2
Table 3. Experimental Mole Fraction Solubility (x) of Compounds Studied in 1-Octanol at Different Temperatures (p = 0.1 MPa)a,b
7.10 8.71 10.58 13.37 16.12 2.07 ± 0.16 3406 ± 48 0.9997 0.9 × 10−2
x × 105
8
3.05 3.77 4.54 6.05 7.36 6.8 ± 0.2 3688 ± 51 0.9997 0.9 × 10−2
x × 103
8
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DOI: 10.1021/acs.jced.7b00641 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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Figure 2. Van’t Hoff plot of mole fraction solubility (lnx) of the compound studied versus reciprocal temperature (T−1) in 1-octanol (a) and hexane (b).
lower than in hexane and 1-octanol and does not exceed 8 × 10−6 mole fraction. In the order of increasing solubility in buffer solution, the compounds are arranged in a series as 1 < 8 < 5 < 3 < 2 < 7 < 6 < 4. Ideal Solubility. The equilibrium solubility of crystalline nonelectrolyte solute in solvent (x) is connected to the ideal solubility (xid) by the following thermodynamic relationship ln x = ln x id − ln γ
This equation is valid on the assumption that the difference in the molar heat capacity of the solid and the supercooled liquid form of the solute at the solution temperature is approximated to be equal to the fusion entropy at the melting point. It should be noted that the ideal solubility calculated from eq 6 is determined only by the properties of individual substances and does not take into account the solvent properties. The obtained values of the ideal solubility (Table 5) exceed the experimental ones in all the systems solute−
(3)
where γ is the activity coefficient of the dissolved substance in the solution (the symmetric standard state). To study the effect of the solid state stability on solubility, we determined the melting enthalpies and temperatures, which are given in Table 2. The solubility dependences of the compounds in 1-octanol and hexane on their melting points and, consequently, the stability of the crystalline phase are shown in Figure 3.
Table 5. Ideal Solubility (xid), Equilibrium Mole Fractions (xo, xw), Distribution Coefficients (K, K0) in System 1Octanol/Buffer pH 7.4 and Activity Coefficients (γo) of Compounds Studied in 1-Octanol at 298 K (p = 0.1 MPa)a,b compound
xid
xo ×102
xw × 105
K
K0
γo
1 2 3 4 5 6 7 8
0.136 0.139 0.263 0.195 0.463 0.262 0.233 0.335
15.67 1.39 15.41 2.96 1.32 1.07 0.93 0.38
77.09 3.75 85.63 11.35 9.16 2.62 1.83 0.45
203.28 370.36 179.84 260.79 144.10 408.16 508.52 841.02
942.29 862.95 751.75 922.59 304.78 795.25 900.48 1107.37
4.63 2.33 4.18 3.54 2.11 1.95 1.77 1.32
a
The standard uncertainties are u(T) = 0.15 K, u(p) = 0.003 MPa. The relative standard uncertainties are ur(x) = 0.04, ur(K) = 0.04, ur(γ) = 0.04. b
solvent, which indicates a positive deviation from ideality (γ > 1). And for all the studied bicycles, there is a correlation between the experimental and calculated solubility values in 1octanol (Figure 4). The dependences presented in Figures 3 and 4 can be used to compare the solubility of compounds of close structure with those studied based on thermophysical data which can be easily determined experimentally. Thermodynamics of Solubility. The activity coefficients of the substances studied in saturated solutions of 1-octanol and hexane were determined by a direct method for the distribution of the compound between two immiscible liquid phases of 1-octanol/buffer pH 7.4 and hexane/buffer pH 7.4 at 298.15 K. The ratio of the activities of the substance in the organic and aqueous phases at the limiting dilution (x → 0 in both phases) is equal to the distribution coefficient and is constant
Figure 3. Solubility correlation with melting temperature of the compound studied in 1-octanol (▲) and hexane (●) (R = 0.920 and 0.863, respectively).
On the basis of the thermophysical characteristics, we calculated the ideal solubility of the compounds studied in octanol according to equation15 ln x1id =
−ΔHm ⎡ Tm ⎤ ln⎢ ⎥ ⎣T ⎦ RTm
(4) E
DOI: 10.1021/acs.jced.7b00641 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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Table 6. Dissolution and Solvation Enthalpies (ΔH0sol, ΔH0solv) of the Compounds Studied in 1-Octanol and Hexane at 298.15 K (p = 0.1 MPa)a ΔH0sol/kJ·mol−1
compound
1-Octanol 12.9 ± 0.2 20.6 ± 0.5 10.7 ± 0.2 29.3 ± 0.5 16.5 ± 0.2 29.1 ± 0.5 34.0 ± 0.9 30.7 ± 0.4 Hexane 49.6 ± 0.9 38.6 ± 0.4 26.9 ± 0.6 30.2 ± 0.5 21.6 ± 0.5 29.4 ± 0.8 30.2 ± 0.3 28.3 ± 0.4
1 2 3 4 5 6 7 8
Figure 4. Correlation between experimental and ideal solubilities of compound studied in 1-octanol (R = 0.881).
K0 =
xoγo ao = = const aw x wγw
1 2 3 4 5 6 7 8
(5)
where ao and aw is the activity, xo and xw is the concentration (mole fraction), γo and γw is the activity coefficient of the compound in the organic phase and aqueous phase, respectively (the asymmetric standard state). As it has been estimated earlier,9 the studied compounds have poor solubility (∼10−6 mol.fraction) in buffer solution pH 7.4, therefore it can be assumed that their activity in the aqueous phase is equal to the concentration (γw = 1) and the following equation can be derived K0 =
a
−88.2 −81.3 −97.5 −63.3 −72.5 −120.9 −90.6 −119.8 −51.5 −63.3 −81.3 −62.4 −67.4 −120.7 −94.4 −122.2
The standard uncertainties are u(T) = 0.15 K, u(p) = 0.003 MPa.
which leads to a solubility decrease. Analysis of the thermodynamic parameters of compounds (4 and 5) shows that an increase in the number of electronegative fluorine atoms in the structure of the substituent leads to a decrease in the dissolution enthalpy. The dissolution of compound (7) with an acetyl substituent is accompanied by the highest enthalpy resulting in a relatively low solubility of the bicycle. The structure of compound (8) differs from (1) by the presence of a chlorine atom in the meta-position of the phenyl ring, which leads to a more significant increase in the enthalpy dissolution and to a decrease in solubility by almost 2 orders of magnitude. The thermodynamic dissolution parameters of the compounds (1 - 6) in hexane, as opposed to 1-octanol, are characterized by strong enthalpy which determine their low solubility in this solvent. The lower enthalpy of dissolution of compound (3) with an isopropyl substituent as compared to compound (2) having a smaller ethyl substituent is observed in both solvents. Solvation. Solvation of the molecules of a dissolved compound is a change in the standard thermodynamic functions under the transfer of one mole of a substance from the gas phase to the solvent.16 This process can be represented as a difference between the thermodynamic parameters of dissolution and sublimation by the following equation
xoγo xw
ΔH0solv/kJ·mol−1
(6)
The distribution coefficients (K = xo/xw) were determined experimentally at different concentrations of the test substance in the organic phase and a dependence of K on the compound concentration was plotted. As example, the dependence of distribution coefficients in system 1-octanol/buffer pH 7.4 (K) versus concentration in 1-octanol phase for compound 1 presented in Figure S2. The values of K0 and K are obtained from this dependence by extrapolating to an infinitely dilute solution (xo = 0) and the concentration of the saturated solution of the substance in an organic solvent (xo are given in Tables 3 and 4). When the distribution in the hexane/buffer pH 7.4 system was studied, it was found that the distribution coefficients do not depend on the compound concentration in hexane. Therefore, the activity coefficients of the substances in hexane solutions were taken equal to unity. The distribution coefficients of the substances in the 1-octanol/buffer pH 7.4 system at an infinite dilution (K0) and at a concentration corresponding to the saturation of the organic solvent (K) as well as the activity coefficients of the substances in 1-octanol (γo) are shown in Table 5. The values of the activity coefficients greater than 1 indicate that thermodynamic forces not favorable the dissolution process of the solute. The thermodynamic dissolution enthalpies of the substances in the solvents studied calculated by eq 1 are summarized in Table 6. Looking at the thermodynamic properties in 1-octanol, it is possible to observe that methyl- (1) and isopropyl- (3) derivatives present the smallest enthalpy of dissolution. The elongation of the alkyl chain meaning the transition from the bicycle with a methyl substituent (1) to the one with an ethyl substituent (2) is characterized by an increase in enthalpy
0 0 0 ΔHsolv = ΔHsol − ΔHsub
ΔH0solv,
ΔH0sol,
(7)
ΔH0sub
where and are the enthalpies of solvation, dissolution, and sublimation. Since the thermodynamic solvation parameters cannot be determined directly from the experiment, two other processes must be studied to obtain this data: dissolution and sublimation. The saturated vapor pressure of the studied crystalline compounds at different temperatures was measured by us previously by the transparent method and the values of the thermodynamic functions of sublimation are given in Table 2. On the basis of these data and the results of the solubility study carried out herein, the solvation enthalpies of the compounds in solutions of 1-octanol and hexane were calculated from eq 7 and summarized in Table 6. In all the F
DOI: 10.1021/acs.jced.7b00641 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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Figure 5. Enthalpies of the compounds studied for the following processes: sublimation (cyan), solvation (gray), and dissolution (red) in 1-octanol (a) and hexane (b).
solute−solvent systems studied herein, the values ΔH0solvare negative. In addition, the solvation processes are enthalpydriven and are accompanied by a decrease in molecular disordering when molecules pass from the vapor to the liquid phase. It is well-known that the energy of crystal lattices has a significant effect on the solubility of substances17 and, as the experiment showed, the test compounds validate this rule. Figures 5a and 5b show histograms of the dissolution, sublimation and solvation enthalpies values of the substances studied in 1-octanol and hexane. We have found that compounds (2−5) with the sublimation enthalpy varying from 89.0 to 108.2 kJ·mol−1 have a higher solubility in both solvents than compounds (6−8) with higher crystal lattice energy varying from 124.6 to 150.5 kJ·mol−1. The sublimation enthalpy of compound (1) was revealed to be rather high, at the same time, it has the maximal solubility in both solvents under study. For the studied bicyclic derivatives, there is a tendency of solvation enthalpy decrease with crystal lattice energy increase (Figure 6). The revealed correlation between
measured by the shake-flask method in two model solvents: 1octanol and hexane. The studied compounds are arranged in the order of decreasing solubility in both solvents at a temperature of 298.15 K in accordance with the structure of the substituent in the para-position of the phenyl ring as follows: methyl-, fluoro-, ethyl-, trifluoromethyl-, cyano- and acetyl-. The solubility of the substances in 1-octanol is approximately 2 orders of magnitude higher than in hexane, and in the studied temperature range T = 293.15 to 315.15 K it varies within 3.0 × 10−3 to 2.1 × 10−1 and 7.1 × 10−5 to 13.5 × 10−3 mole fraction for solutions of alcohol and alkane, respectively. The results are explained by the formation of a hydrogen bond between the 1-octanol molecule, which acts as a proton donor, and the tertiary nitrogen atom of the heterocycle, which is an electron donor center in the structure of the compounds under study. In all the studied solute− solvent systems, the experimental values of solubility in 1octanol are correlated with the melting temperatures of the compounds, as well as with the values of ideal solubility calculated based on the thermophysical characteristics. By the method of substance distribution between two liquid phases, the activity coefficients of the compounds in the saturated solutions of the used solvents were determined. The dissolution enthalpies were calculated on the basis of solubility temperature dependences. The positive ΔH0sol revealed that dissolution in each solvent was an endothermic process. It has been found that the dissolution enthalpies in hexane are characterized by a strong values and in contrast to 1-octanol, which leads to low solubility of the substances in this solvent. On the basis of the results of the investigation into the dissolution and sublimation processes of the studied bicyclic derivatives, we calculated the thermodynamic solvation enthalpies. A tendency of solvation enthalpy reduction in both solvents with an increase in the crystal lattice energy of the substances has been revealed.
Figure 6. Correlation between solvation and sublimation enthalpies of the compound studied in 1-octanol (●) and hexane (▲) for compounds studied (R = 0.9171).
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the sublimation and solvation enthalpies indicates that the substances with stronger intermolecular interaction in the crystalline state are better solvated.
ASSOCIATED CONTENT
S Supporting Information *
The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.7b00641.
4. CONCLUSION The solubility of eight biologically active nonaromatic compounds with a common 3-thia-1-aza-bicyclo fragment and substituents of different chemical nature in the phenyl ring was
1
H NMR spectra of compounds studied and concentration dependence of distribution coefficients for compound 1 in system 1-octanol/buffer pH 7.4 (PDF)
G
DOI: 10.1021/acs.jced.7b00641 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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AUTHOR INFORMATION
Corresponding Author
*Phone: 7 (4932)351545. Fax: 7(4932)336246. E-mail: avs@ isc-ras.ru. ORCID
Angelica V. Sharapova: 0000-0002-5841-2724 Tatyana V. Volkova: 0000-0002-0562-6999 German L. Perlovich: 0000-0002-6267-5244 Notes
The authors declare no competing financial interest.
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REFERENCES
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DOI: 10.1021/acs.jced.7b00641 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
