Article pubs.acs.org/jced
Thermodynamics of 1,3-Dimethylurea in Eight Alcohols Peipei Zhu,† Yanxin Chen,‡ Yanan Zhou,† Yan Xiao,† Jinbo Ouyang,† Xin Huang,† Gaohong He,§ Baohong Hou,†,∥ and Wei Chen*,†,∥ †
National Engineering Research Center of Industry Crystallization Technology, School of Chemical Engineering and Technology, Tianjin University, Tianjin 300072, China ‡ Sinopec Shanghai Research Institute of Petrochemical Technology, Shanghai 201208, China § State Key Laboratory of Fine Chemicals, R&D Center of Membrane Science and Technology, School of Chemical Engineering, Dalian University of Technology, Dalian 116024, China ∥ Collaborative Innovation Center of Chemical Science and Chemical Engineering (Tianjin), Tianjin 300072, China S Supporting Information *
ABSTRACT: In this work, solubility data of 1,3-dimethylurea (DMU) in methanol, ethanol, 1-propanol, 1-butanol, 1-pentanol, 1-hexanol, 1-heptanol, and 1-octanol were measured by synthetic method from T = 288.15 to 328.15 K at ordinary pressure. It turned out that the solubility data increased as temperature and solvent polarity increased. The molecular interactions in solution have been studied by the linear energy relationship. Furthermore, based on two thermodynamic models (i.e., modified Apelblat and Wilson), the experimental results were fitted and analyzed. Finally, the thermodynamic properties of solution, including entropy, enthalpy, and Gibbs energy were calculated. These results indicate that in all selected solvents the dissolution behavior were endothermic, entropy-driven, and not spontaneous.
1. INTRODUCTION 1,3-Dimethylurea (DMU, C3H8N2O, CAS No. 96-31-1), shown in Figure 1, is regarded as the drug intermediate (including
data were obtained successfully. According to the different thermodynamic properties of pure component (including melting temperature, enthalpy of fusion, mole volume, and so on), experimental data were fitted and analyzed with three models (i.e., linear energy relationship, modified Apelblat and Wilson equations). Finally, the dissolution entropy together with the enthalpy and Gibbs free energy change during DMU in above-mentioned solvents were determined according to the van’t Hoff model.
Figure 1. Molecule structure of DMU.
2. EXPERIMENTAL SECTION
theophylline, caffeine, and so on).1 The purity of DMU is a mandatory requirement for its pharmaceutical application and usually improved by crystallization processes. Since in crystallization the supersaturation of solution is crucial to the quality of the crystals, the fundamental solubility data of DMU are essential factors in development, design, and control of crystallization processes. Accordingly, it is vital to investigate solubility behavior of DMU in different commercial solvents to ensure the manufacturing process with desired yield and purity can be designed.2 In addition, thermodynamic properties of DMU’s dissolution in various solvents can also be deduced from corresponding solubility data. Hence, this work is used to obtain the solubility results of DMU in methanol, ethanol, 1-propanol, 1-butanol, 1-pentanol, 1-hexanol, 1-heptanol, and 1-octanol by a synthetic method between T = 288.15 and 328.15 K at ordinary pressure. By using the laser monitoring observation technique the solubility © XXXX American Chemical Society
2.1. Materials. The white crystal powder of DMU with the purity more than 99% was offered by Aladdin Industrial Corporation, Shanghai, China and stored in a desiccator used without any other treatment and identified as form I. The experimental solvents (methanol, ethanol, 1-propanol, 1butanol, 1-pentanol, 1-hexanol, 1-heptanol, and 1-octanol) were bought from Tianjin Jiangtian Chemical Technology Co., Ltd., China and of analytical reagent (AR). Before experiments, the molecular sieves were used to remove the original water in the selected solvents. The details, including purities and sources of all of the materials used in this work, are listed in Table 1. Received: October 21, 2015 Accepted: March 16, 2016
A
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Table 1. Sources and Mass Fraction Purity of the Materials chemical name DMU methanol ethanol 1-propanol 1-butanol 1-pentanol 1-hexanol 1-heptanol 1-octanol a
source Aladdin Industrial Co., Shanghai, China Jiangtian Chemical Technology Co., Ltd., Jiangtian Chemical Technology Co., Ltd., Jiangtian Chemical Technology Co., Ltd., Jiangtian Chemical Technology Co., Ltd., Jiangtian Chemical Technology Co., Ltd., Jiangtian Chemical Technology Co., Ltd., Jiangtian Chemical Technology Co., Ltd., Jiangtian Chemical Technology Co., Ltd.,
Tianjin, Tianjin, Tianjin, Tianjin, Tianjin, Tianjin, Tianjin, Tianjin,
mass purity
purification method
analysis method
0.998 >0.995 >0.995 >0.996 >0.995 >0.995 >0.995 >0.995 >0.995
none none none none none none none none none
HPLCa GCb GCb GCb GCb GCb GCb GCb GCb
China China China China China China China China
High-performance liquid chromatography. bGas−liquid chromatography.
2.2. Procedure of Differential Scanning Calorimetry. By using indium and zinc standards, the Mettler-Toledo DSC 1/500 was calibrated, and then the melting temperature Tm and enthalpy of fusion ΔfusH of DMU were determined. About 5 mg of DMU powder was put into a hermetic DSC pan which was measured by the electronic analytical balance (MettlerToledo AE240). At the nitrogen atmosphere (2.5 mL·s−1), this sample was scanned uniformly from 303.15 to 403.15 K with a constant heating rate (2 K·min−1). The same experiment was repeated three times. 2.3. Solubility Measurement. With a synthetic method,3 the solubility data of DMU in different solvents were measured. From the exiting literature,4 the details about the measuring method can be obtained. It contained a laser beam and a 100 mL jacketed glass vessel. In order to provide a constant temperature and stirring, the external thermostat (CF41, Julabo, Seelbach, Germany) and the magnetic stirring bar were used. The standard uncertainty of system temperature was 0.03 K. During the measurement, the electronic analytical balance Mettler-Toledo (AB204-N) was used. First, an insufficience of solvent as well as excess DMU crystals were added into the jacketed vessel. Next, the solution was kept continuously stirring until the change of the temperature less than 0.05 K. At last, the selected solvent into the vessel was added slowly until the final solid solute disappeared, and at this time, through the vessel, the intensity of the laser reached the maximum. So all amounts of the solute and solvent can be calculated in this experiment. All of the experiments were repeated three times in order to verify the uncertainties. The mole fraction solubility x1 can be obtained by the following equation: x1 =
m1/M1 m1/M1 + m2 /M 2
liquid equilibrium behavior of DMU in all selected solvents can be explained accurately. 3.1. Modified Apelblat Equation. Modified Apelblat equation, as widely known, was derived from the Clausius− Clapeyron equation. As previously reported,5 it could be applied in explaining the relationships between solubility and temperature. So the concrete equation can be described as5 ln x1 = A +
B + C ln T T
(2)
where A, B, and C stand for the empirical constants obtained by fitting experimental solubility data and T is the absolute temperature. 3.2. Wilson Model. According to the Wilson equation6 in the binary system, the nonideality of mixing process could be described. The following equation can be obtained on the basis of the Wilson model: ⎞ ⎛ Λ12 Λ 21 ln γ1 = −ln(x1 + Λ12x 2) + x 2⎜ − ⎟ x 2 + Λ 21x1 ⎠ ⎝ x1 + Λ12x 2 (3)
where Λ12 =
⎡ λ − λ11 ⎤ V2 exp⎢ − 12 ⎥ ⎣ V1 RT ⎦
(4)
Λ12 =
⎡ λ − λ 22 ⎤ V1 exp⎢ − 21 ⎥ ⎣ V2 RT ⎦
(5)
Λ12 and Λ21 are Wilson parameters and V1 and V2 represent the molar volumes of solute and solvent, respectively. As the tunable parameters of Wilson model, Δλ12 = (λ12 − λ11) and Δλ21 = (λ21 − λ22) can be used to indicate the interaction energy.
(1)
4. RESULTS AND DISCUSSION 4.1. Form Identification of DMU. Compared to the corresponding theoretical patterns from the literature,7 the commercial available DMU can be identified as form I as shown in Figure S1 of the Supporting Information, and there is no obvious polymorphic transition peak in Figure 2. Both indicate the sample DMU contains no detectable form II phase. Although it was mentioned DMU exhibited a solid-phase transition at 301.2 K (form II to I) in the reference,8 the transition was not observed in this work. Also, since solid DMU were added in the clear solution in small amounts many times during a synthetic process (observable suspension meaning saturation), the effect of phase transition on solubility was minimized during measurements.
where m and M stand for the total masses used in this experiment and molar masses, and subscript 1 and 2 represent the solute and solvent, respectively. 2.4. X-ray Powder Diffraction Analysis. By using the Xray powder diffraction (XRPD, Rigaku D/max-2500, Rigaku, Japan), the crystal form of DMU was measured. It is performed with the Cu Kα radiation (0.071073 Å) in the 2-theta range from 2° to 50°a with a scanning rate of 0.067 deg·s−1, under 100 mA current and 40 kV voltage.
3. THERMODYNAMIC MODELS The experimental data were fitted and analyzed with two models (i.e., modified Apelblat and Wilson), so that the solid− B
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Table 2. Experimental and Calculated Mole Fraction Solubility of DMU in the Selected Alcohols with the Temperature Range from 288.15 to 328.15 K under p = 101.3 kPaa
Figure 2. DSC scan of DMU as shown in ref 9.
4.2. Physical Properties of Pure Materials. As shown in Figure 2 and ref 9, the result of the mean extrapolated onset temperature is Tm = 376.35 K. This value is in good accordance with the existing data, such as Barone et al.10 (Tm = 375.15 K), Parnham et al.11 (Tm = 374.15−378.15 K), Ferro et al.12 (Tm = 379.4 K), Gatta et al.13 (Tm = 379.46 K), and Zordxa et al.14 (Tm = 379.87 K). The molar enthalpy of fusion ΔfusH = 12.76 kJ·mol−1 shows a good agreement with the literature values of ΔfusH = 12.76 kJ·mol−1.15 There still exist other relevant values; for example, Gatta et al.13 obtained the value of ΔfusH = 13.62 kJ·mol−1, while Zordxa et al.14 reported 12.64 kJ·mol−1. These slight deviations in the results may be from the differences of measurement method, sample source, experiment environment or other related factors. The standard uncertainty for the melting temperature was 0.3 K, and the relative standard uncertainty for enthalpy of fusion was 0.01. All physical properties of the used pure materials are showed Table S1 of the Supporting Information including the molar mass and the density of the solvents used in this work and reported in literature.16−21 With a pycnometer method,22 the densities of the selected solvents in this work were obtained. 4.3. Solubility Analysis. From Figure 3 and Table 2, the experimental and fitting solubility value of DMU in methanol,
T/K
x1exp
288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15
0.384 0.413 0.448 0.490 0.530 0.571 0.627 0.674 0.719
288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15
0.336 0.359 0.396 0.438 0.468 0.526 0.570 0.621 0.666
288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15
0.317 0.339 0.373 0.416 0.463 0.498 0.550 0.590 0.629
288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15
0.308 0.328 0.369 0.407 0.434 0.471 0.518 0.568 0.615
x1Apel Methanol 0.380 0.415 0.451 0.490 0.531 0.575 0.621 0.671 0.723 Ethanol 0.331 0.363 0.398 0.436 0.477 0.521 0.567 0.617 0.671 1-Propanol 0.308 0.343 0.380 0.419 0.459 0.500 0.543 0.588 0.633 1-Butanol 0.306 0.335 0.366 0.400 0.436 0.476 0.519 0.565 0.615
x1Wilson
x1exp
0.384 0.414 0.448 0.489 0.530 0.572 0.626 0.674 0.720
0.300 0.325 0.361 0.392 0.425 0.465 0.487 0.529 0.568
0.336 0.361 0.396 0.436 0.469 0.523 0.569 0.621 0.670
0.296 0.321 0.358 0.382 0.419 0.457 0.481 0.521 0.560
0.318 0.342 0.374 0.414 0.458 0.496 0.547 0.591 0.635
0.277 0.320 0.345 0.376 0.417 0.444 0.474 0.507 0.544
0.309 0.331 0.367 0.403 0.433 0.471 0.518 0.568 0.617
0.242 0.268 0.303 0.331 0.373 0.403 0.444 0.479 0.519
x1Apel 1-Pentanol 0.299 0.328 0.359 0.392 0.425 0.459 0.494 0.530 0.566 1-Hexanol 0.295 0.324 0.354 0.385 0.418 0.452 0.486 0.522 0.559 1-Heptanol 0.282 0.313 0.346 0.379 0.412 0.445 0.477 0.510 0.541 1-Octanol 0.240 0.270 0.301 0.334 0.369 0.405 0.442 0.480 0.519
x1Wilson 0.303 0.328 0.359 0.389 0.422 0.460 0.488 0.530 0.572 0.299 0.323 0.355 0.381 0.416 0.453 0.482 0.522 0.564 0.284 0.318 0.344 0.374 0.411 0.441 0.473 0.509 0.549 0.245 0.270 0.301 0.330 0.368 0.401 0.442 0.480 0.523
a
Standard uncertainties are u(T) = 0.03 K, ur(x1) = 0.02, and u(p) = 0.3 kPa.
ethanol, 1-propanol, 1-butanol, 1-pentanol, 1-hexanol, 1heptanol, and 1-octanol can be obtained. Figure 3 shows a clear tendency that with the increase of temperature the solubility values of DMU increased in the various alcohols. Compared with other solvents selected in this work, the solubility value of DMU in 1-octanol shows the lowest while in methanol is the highest one throughout the whole temperature range. The order of the solubility data in the selected alcohols is shown as follows: 1-octanol < 1-heptanol < 1-hexanol < 1pentanol < 1-butanol < 1-propanol < ethanol < methanol. Research indicates that this sequence is in good accordance with the polarity order23 and Hildebrand solubility parameter24 order of the solvents which is shown in Table S2 of the Supporting Information. More importantly, further study shows
Figure 3. Mole fraction solubility x1 versus T in the selected solvents: (●) methanol; (○) ethanol; (▲) 1-propanol; (△) 1-butanol; (■) 1pentanol; (□) 1-hexanol; (★) 1-heptanol; (☆) 1-octanol. The corresponding lines are correlation results based on the Wilson model. C
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Table 3. Optimized Model Parameters and Deviations in Different Alcohols at p = 101.3 kPaa solvents models modified Apelbat
Wilson
a
A B C 102 RAD 102 RMSD Δλ12 Δλ21 102 RAD 102 RMSD
methanol
ethanol
1-propanol
1-butanol
1-pentanol
1-hexanol
1-heptanol
1-octanol
−21.40 −343.5 3.819 0.5372 0.3426 −836.0 5855 0.1584 0.1011
−27.90 −183.7 4.844 0.9599 0.4823 352.6 5353 0.3186 0.1977
48.02 −3678 −6.432 1.157 0.529 1044 4953 0.5676 0.3157
−51.57 910.5 8.338 0.8671 0.4244 1505 4747 0.3905 0.2045
41.74 −3233 −5.602 0.5720 0.3364 1711 4975 0.6702 0.2956
30.65 −2725 −3.957 0.6777 0.3183 2073 4747 0.5803 0.2627
105.7 −6182 −15.10 0.8880 0.3868 2471 4494 0.8548 0.3949
65.97 −4610 −9.075 0.5660 0.2231 3364 3918 0.6780 0.2671
The standard uncertainty is u(p) = 0.3 kPa.
that, with the decrease of the number of CH2 units, the solubility data increased. According to the linear solvation energy relationship,25 the chemical properties of solute and solvent, such as the possibility and ability to form hydrogen bonds, the polarity and structure can be related to the solution thermodynamic properties, solution property, different distribution cases, enthalpy, entropy, free energy in a state of equilibrium, and so on.25 Researches indicate that the solvent−solvent, solute−solute, and solvent−solute interactions can influence the process of solute dissolution, as well as the thermodynamic properties of solvent. A multiple linear regression analysis (MLRA) has been found in order to get a deep insight into these thermodynamic properties. This analysis method involves a serious of solvent parameters. According to this, in the solute−solvent system, the Gibbs free energy (XYZ) can be written as a linear relationship with various parameters. The specific equation is described as follows:26
log x1 = − 2.2312 + 1.2256α + 0.4663β + 1.7848π * − 0.7413
R2 = 0.955, (7)
As shown above, the positive values of π*, α, and β indicate that, as the values of these parameters increase, the solubility data increase. It could be observed that the ability of hydrogen bond acceptance (HBA) of carbonyl group of DMU with hydroxyl group of alcohol, the ability of hydrogen bond donation (HBD) of the methylamino groups with hydroxyl group, and polar interactions between solute and solvent favor the increase of DMU’s solubility. It can be seen that the value of the δH is negative which indicates that with the decreasing of the self-cohesiveness (or structuredness) the solubility in the solvent increases.26 Hence, the self-cohesiveness related to the solvent−solvent interaction show negative effects on the solubility of DMU. Above all, this parameter of π* shows a most important weight in DMU solubility, while others also exert considerable influence. On the other hand, the magnitude of the coefficients of β is smaller compared to those of other parameters indicating that the DMU solubility is less susceptible to the variations of β. In order to further describe the solid−liquid equilibrium more quantitatively, the temperature and solubility values were correlated by two thermodynamic models (modified Apelblat and Wilson). Via fitting the experimental data the parameters of these models were obtained with each thermodynamic model. The calculated solubilities of the DMU in the selected alcohols are shown in Table 2 together with the experimental data. The accuracy of these model correlation was evaluated by using the root-mean-square deviations (RMSDs), together with the relative average deviation (RAD), which can be obtained by the following equation:
∑ solute−solvent interaction energy
Among them, the XYZ is the parameter that showed a linear relationship with the Gibbs free energy, while XYZ0 represents the constant which is related to the solute. Moreover, this equation can be extended to all the forms of solute−solvent interaction. According to Kamlet−Taftthe, as described in the existing literature, it can be written as follows:24 XYZ = XYZ0 + aα + bβ + pπ * + mδ H2
n = 12, T = 298.15 K
σ = 0.0006,
XYZ = XYZ0 + cavity formation energy +
δ H2 1000
(6)
where m, p, a, and b are the coefficients of the linear regression which are related to the solvent. These values changed along with the solvent properties. α and β are the solvent parameters about the solvatochromic properties. Among them, α represent the ability of hydrogen bond donation (HBD), while β is the ability of hydrogen bond acceptance (HBA). The π* is the parameter of polarity and polarizability, while δH2 show the Hildebrand solubility parameter which is on behalf of the selfcohesiveness of the solvent. All of the parameters show a good relationship with the solvent. The values of α, β, π*, and δH2 of the selected solvents which are used in this work are shown in Table S2 of the Supporting Information, and according to these, the results for solubility can be expressed as follows:
N
RMSD =
RAD =
1 N
∑i = 1 (xiexp − xical)2
N
∑ i=1
N
(8)
xiexp − xical xiexp
(9)
where N represents the number of experimental points and xcal i and xexp stand for calculated and experimental data of DMU, i respectively. In Table 3, the different parameters of the selected models can be obtained together with the RMSD and RAD. From the D
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where N stands for the number of temperatures studied. In current work, Tmean is calculated to be 307.61 K. All of the thermodynamic properties could be seen as constant when the temperature changes a little. Because of this, the heat capacity of solution could also be considered as invariant. Thus, the apparent standard enthalpy (ΔHsol) and molar Gibbs free energy (ΔGsol) change could be determined from the following equations based on van’t Hoff analysis:30
very small RMSDs, it can be deduced that the calculated data of DMU are in good accordance with the values of this work in all pure alcohols. For example, the relative average deviations (RADs) of the solubility data of DMU in alcohols correlated with Wilson model are as follows: 0.12%, 0.29%, 0.52%, 0.36%, 0.66%, 0.57%, 0.84%, and 0.66%, respectively. It indicates that all of the absolute values of these are all less than 0.9%. Thus, the Wilson model can be considered as the suitable model in fitting the solubility data of DMU in all the selected alcohols. Besides, the fitting values of the modified Apelblat equation indicate that it could also be the suitable model to fit. However, comparing with the modified Apelblat, the Wilson model shows the best fitting values. The correlation lines with the solubility data x1 versus T in the solvents are plotted in Figure 3 which are obtained from the Wilson model. As shown in Figure 3 and 4, although temperature dependences of solubilities in eight
⎛ ⎞ ⎡ ∂ ln x1 ⎤ ∂ ln x1 ΔHsol = −R ⎢ ⎟ ⎥ = −R ⎜ ⎣ ∂(1/T ) ⎦ ⎝ ∂(1/T − 1/Tmean) ⎠
(11)
ΔGsol = −RTmean × intercept
(12)
The ln x versus 10 (1/T − 1/Tmean) plots are listed in Figure 4; by using this figure, the ΔGsol and ΔHsol could be obtained from the intercept and slope of the line, respectively. The slope, intercept, and R2 of each plot are shown in Table 4. All of the 4
Table 4. Slope and Intercept of the lnx1 vs 104(1/T − 1/ Tmean) Plot in the Pure Solvents slope
R2
methanol ethanol 1-propanol 1-butanol 1-pentanol 1-hexanol 1-heptanol 1-octanol
−0.6381 −0.7459 −0.7948 −0.8319 −0.8698 −0.8846 −0.9100 −1.0158
−0.1515 −0.1668 −0.1694 −0.1653 −0.1513 −0.1510 −0.1546 −0.1820
0.9985 0.9967 0.9965 0.9965 0.9979 0.9981 0.9923 0.9985
ΔSsol =
solvents show a similar tendency and all of data are within the permissible error range, those in 1-pentanol, 1-hexanol, and 1heptanol slightly deviate from those observed in other alcohols. The different solubility behavior could be a result from multiply factors, such as the quality of reagents, humidity of laboratory environment, and experimental errors. Despite similar deviations also could be found in solubilities of some compounds (for example lovastatin) in series of alcohols,27 few papers discussed the phenomena. Accordingly, in the future, more research will be needed to reveal the actual reasons. 4.4. Thermodynamic Functions of Solution. As known widely, researching the behavior of the solute dissolution in kinds of solvents are of significant for the real solution. In the current work, the thermodynamic properties of solution, including entropy, enthalpy and Gibbs energy were determined based on the experimental data. According to the van’t Hoff equation, the entropy and enthalpy of solution can be deduced from the plot of ln x vs 1/T, and this method could be found widely used28 in existing research.29 As shown in the literature, the average temperature was determined as the following which was used to reduce the error of the calculation.
ΔHsol − ΔGsol Tmean
(13)
In order to compare the contribution to Gibbs free energy change between enthalpy and entropy, %ζH and %ζTS are deduced. They stand for the contribution values of enthalpy and entropy and are determined in the following equation, respectively:31 %ζH = 100 ×
|ΔHsol| |ΔHsol| + |TmeanΔSsol|
(14)
%ζTS = 100 ×
|TmeanΔSsol| |ΔHsol| + |TmeanΔSsol|
(15)
The calculated results of eqs 11−15 are shown in Table 5. As seen in Table 5, it could be obtained that all of the apparent standard Gibbs energy of solution are positive, indicates that the process of dissolution is nonspontaneous obviously, which is in good agreement with existing literature.32,33 This result may be attributed to the fact that, where using the solid pure solute as the reference state having unit as solute concentration, the concentration scale (mole fraction) used for most nonelectrolyte compound is much lower than that. Because of this its logarithmic term is obviously negative; hence, ΔGsol obtains a positive quantity. Table 5 also
N i=1
intercept
R2 are more than 0.9923, some even up to 0.9985. It reflects that it is suitable to use the modified van’t Hoff equation to determine the thermodynamic properties of DMU solution. Then using the Gibbs equation, the standard apparent entropy change of solution (ΔSsol) at average temperature can be determined as follows:30
Figure 4. Modified van’t Hoff plot of the mole fraction solubility of DMU in selected solvents: (★) methanol; (☆) ethanol; (▲) 1propanol; (△) 1-butanol; (●) 1-pentanol; (○) 1-hexanol; (■) 1heptanol; (□) 1-octanol.
Tmean = N /∑ (1/Ti )
solvent
(10) E
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Table 5. Thermodynamic Functions Relative to the Solution Process of DMU in the Selected Solvents at the Mean Armonic Temperature (Tmean = 307.61 K)a
a
solvent
ΔGsol (J·mol−1)
ΔHsol (J·mol−1)
ΔSsol (J·mol−1·K−1)
TmeanΔSsol (kJ·mol−1)
%ζH
%ζTS
methanol ethanol 1-propanol 1-butanol 1-pentanol 1-hexanol 1-heptanol 1-octanol
1632 1908 2033 2128 2225 2262 2327 2598
12596 13872 14086 13743 12576 12553 12855 15134
35.64 38.89 39.18 37.76 33.65 33.45 34.22 40.75
10.96 11.96 12.05 11.62 10.35 10.29 10.53 12.54
53.46 53.69 53.89 54.20 54.85 54.95 54.98 54.69
46.54 46.31 46.11 45.80 45.15 45.05 45.02 45.31
Combined expanded uncertainties U are Uc(ΔGsol) = 0.065ΔGsol, Uc(ΔHsol) = 0.055 ΔHsol, Uc(ΔSsol) = 0.060 ΔSsol (0.95 level of confidence).
China (21376164 and 21376165), and Tianjin Municipal Natural Science Foundation, China (No. 13JCZDJC28400).
indicates that in the studied cases the solubility of DMU apparently increases with decreasing standard Gibbs energy of solution. All of the above data show that the results of ΔGsol increases as the DMU solubility decreases. It can be found that the ΔHsol and ΔSsol of DMU in all selected solvents are all positive, which deduced that the process of dissolving is apparently not only endothermic but also entropy-driving, which gives a good reason to explain the increasing DMU solubility with increasing temperature. Moreover, the enthalpy (%ζH greater than 53.4) is the main contribution to the molar Gibbs free energy change of DMU of all cases.
Notes
The authors declare no competing financial interest.
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5. CONCLUSIONS The solubility values of DMU in eight alcohols, including methanol, ethanol, 1-propanol, 1-butanol, 1-pentanol, 1hexanol, 1-heptanol, and 1-octanol, were measured by the synthetic method from T = 288.15 to 328.15 K at ordinary pressure. It was found that, with the increase of temperature and solvent polarity, the solubility increased. All of the data were correlated and analyzed with the linear energy relationship and modified Apelblat and Wilson models. According to the linear Gibbs free energy relationship, the solubility increased with hydrogen bonding and polar interactions between DMU and solvent while decreasing with the solvent’s self-cohesiveness. Furthermore, enthalpy, entropy, and Gibbs free energy of solution in the selected alcohols were determined based on the van’t Hoff model. The positive values indicated that the process of dissolution in all selected alcohols were endothermic and entropy-driven.
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ASSOCIATED CONTENT
* Supporting Information S
The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.5b00882. Some physicochemical properties of DMU and the selected solvents; results of solubility of DMU in different pure solvents and parameters of solvents at 298.15 K; X-ray powder diffraction pattern (PDF)
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The authors are grateful for the financial support from the Major National Scientific Instrument Development Project of China (21527812), National Natural Science Foundation of F
DOI: 10.1021/acs.jced.5b00882 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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DOI: 10.1021/acs.jced.5b00882 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
