Article pubs.acs.org/jced
Solubility of L-Phenylalanine Anhydrous and Monohydrate Forms: Experimental Measurements and Predictions Jie Lu,*,† Qing Lin,† Zhen Li,† and Sohrab Rohani‡ †
National Engineering Laboratory for Cereal Fermentation Technology, School of Chemical & Material Engineering, Jiangnan University, Wuxi 214122, China ‡ Department of Chemical and Biochemical Engineering, The University of Western Ontario, London, Ontario N6A 5B9, Canada
ABSTRACT: In this work the solubility of L-phenylalanine anhydrous and monohydrate forms in pure water, a water + acetone mixture, and a water + ethanol mixture at various temperatures were measured systematically using the gravimetric method as well as predicted using the universal quasichemical (UNIQUAC) model. Generally the solubility increased with the temperature and decreased with an increase in the content of acetone or ethanol in the solvent mixtures. The two forms were found to be enantiotropically related, and the transition temperature was about 308.9 K. Besides, the UNIQUAC model was demonstrated to work well for the prediction of the solubility of this pseudopolymorphic system in water and its binary mixtures with a low content of organic solvent.
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models.6 EOS's are widely used for phase equilibrium calculations; however, predicting the solid−liquid equilibria of nonideal systems using EOS with ordinary parameters is generally not reliable, and special mixing rules are required to use such equations.7 On the other hand, solubility prediction using activity coefficient models is based on the estimation of the solute ideal solubility and activity coefficient of the solute in a particular solvent.8 Using the thermal properties of a solute, for example, the heat of fusion and melting temperature, the ideal solubility of the solute can be calculated. Then activity coefficients will be calculated from a particular model, which expresses the excess Gibbs free energy of the mixtures as a function of the composition.9−11 To date, various semipredictive models such as Wilson,12 nonrandom two-liquid (NRTL), and modified models,13−17 universal quasichemical (UNIQUAC) and modified models,18−20 and so forth, are widely used to calculate the activity coefficients due to the requirement for only a small number of adjustable parameters to be known, together with structural parameters that may be readily calculated or read from tables.21,23
INTRODUCTION The well-controlled crystallization of active pharmaceutical ingredients (APIs) is often vitally important to the pharmaceutical industry. This is because crystallization is not only in general the last chemical purification step in the production of APIs but also an effective means to control the quality of APIs in terms of crystal form, shape, size distribution, bulk and tap density, filterability, and flowability, which have the potential to impact bioperformance.1 The solubility is a prerequisite to determine the throughput, yield, and driving force of crystallization operations. Thus, knowing the solubility is essential in the study of a crystallization process.2,3 However, the experimental determination of the solubility of pharmaceuticals in various solvents requires a large amount of the solute and is usually a time-consuming procedure. In many cases, there are only a few milligrams of an expensive pharmaceutical available to perform a large number of solubility measurements; therefore solubility prediction over a wide range of temperatures using thermodynamic models is a preferred choice.4 Solubility prediction is very difficult because it depends on a large number of factors such as the chemical structure, solute− solvent interactions, and temperature.5 Traditionally there are two classes of thermodynamic models for phase equilibrium calculations: equations of state (EOS) and activity coefficient © 2012 American Chemical Society
Received: January 2, 2012 Accepted: April 5, 2012 Published: April 15, 2012 1492
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The characterization of two forms was conducted by differential scanning calorimetry (DSC), thermogravimetric analysis (TGA), and FTIR. FTIR spectra were recorded from KBr disks using a FTLA2000-104 spectrophotometer (ABB Bomem, Canada). Ground KBr powder was used as the background in the measurements. The number of scans was 32, and the resolution was 4 cm−1. The measured wavenumber range was from (4000 to 400) cm−1. DSC was performed using a Mettler-Toledo DSC-822e differential scanning calorimeter (Mettler-Toledo, Columbus, OH). Indium was used for calibration. Accurately weighed samples ((5 to 8) mg) were placed in hermetically sealed aluminum pans and scanned at 10 K·min−1 under a nitrogen purge. TGA was conducted with a Mettler-Toledo TGA/DSC 1/1100 SF instrument (MettlerToledo, Columbus, OH) that was also calibrated with indium prior to analysis. The powders of the each form were heated at 10 K·min−1 to 623.15 K under nitrogen purge.
For strongly nonideal systems, the UNIQUAC equation often provides a good representation of equilibrium.24 In this work, the solubility of the anhydrate and monohydrate of L-phenylalanine which is widely used in pharmaceutical and food industries25 are measured in pure water, a water + ethanol mixture, and a water + acetone mixture and predicted using the UNIQUAC model.
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EXPERIMENTAL SECTION Mohan et al.26 indicated that L-phenylalanine (IUPAC name: 2-amino-3-phenylpropanoic acid) can exist in two different crystalline states, anhydrous form and monohydrate form. Anhydrous L-phenylalanine (abbreviated to L-phe) was purchased from Sigma-Aldrich (St. Louis, MO). Table 1 lists Table 1. Purities and Suppliers of Chemical and Solvents
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molar mass chemical name L-phenylalanine
ethanol acetone a
supplier
mass fraction purity
g·mol−1
Sigma-Aldrich Sigma-Aldrich Sigma-Aldrich
≥ 0.980 ≥ 0.998b ≥ 0.998b
165.19 46.07 58.08
a
SOLID−LIQUID EQUILIBRIUM For nonideal solutions the relation between solid and liquid at equilibrium is given as8 ⎛ f L ⎞ ΔH ⎛ T ⎞ ΔCP ⎛ Ttp ⎞ tp fus ln⎜⎜ 2S ⎟⎟ = − 1⎟ − − 1⎟ ⎜ ⎜ RTtp ⎝ T R ⎝T ⎠ ⎠ ⎝ f2 ⎠ ΔCP Ttp + ln R T
Biochemical reagent. bGas chromatography.
the specifications and suppliers of the chemical and solvents used in this work. All of them were used without further purification. Ultrafiltered, deionized water was obtained with a Direct-Q Millipore system (Millipore, Billerica, MA). The L-phenylalanine monohydrate (abbreviated to L-phe·H2O) was obtained by recrystallization: A saturated solution of anhydrate was prepared in water at 343.15 K and then quenched in a thermostat maintained at 278.15 K. After stirring for 24 h, crystals of the monohydrate formed and then were filtered and dried for 24 h at 303.15 K under vacuum. The solubilities of two forms were measured in water + acetone and water + ethanol mixtures with different mass fractions at various temperatures by a gravimetric method: Excess amounts of each form were added to the solvent in a 100 mL jacketed glass crystallizer of which the temperature was controlled by a Julabo 32 ME programmable circulator with an uncertainty of ± 0.01 K (Julabo, Seelbach, Germany). It was found experimentally that about 20 min of proper stirring was enough for the dissolution of solid forms to very closely approach to saturation; meanwhile, the induction time for the transformation from anhydrous form to monohydrate form generally took more than 2 h in the temperature range studied in this work. Thus after 1 h, the agitation was stopped, and the solution was allowed to settle for 0.5 h. The supernatant was then filtered through Millex-VV 0.45 μm filters (Millipore, Billerica, MA); meanwhile, the residue was checked by Fourier transform infrared spectroscopy (FTIR). About 5 mL of the filtered supernatant was withdrawn with a pipet and placed in a sample bottle preweighed by use of a JA1203 analytical balance (Hangping, Wuxi, China) with a resolution of ± 0.1 mg. The samples of anhydrous and monohydrate forms were placed in Vacucenter VC 20 ovens (SalvisLab Renggli AG, Sursee, Switzerland) and vacuum-evaporated to dryness at (318.15 and 303.15) K, respectively, until the mass was constant. The solid residue mass was determined, and the concentration was then calculated. All experiments were replicated three times. The data reported in this work are the average of the replicates, and the uncertainties U are calculated by the Bessel method.
(1)
where f 2L and f 2S are pure subcooled liquid and pure solid fugacities, respectively, and Ttp is the triple point temperature which can be considered as the melting point, Tm. ΔHfus is heat of fusion, and ΔCP is the difference in heat capacities of liquid and solid at T. Fugacities can be related to each other by
x 2 γ2 =
f 2S f2L
(2)
where γ2 is the activity coefficient of solute in the solution and x2 is the solubility of solute in real solution in terms of mole fraction. For ideal systems, the activity coefficient of the solute is equal to unity, and the general equation of ideal solubility can be written as ln
1 x 2ideal
=
⎞ ⎞ ΔCP ⎛ Ttp ΔHfus ⎛ Ttp − 1⎟ − − 1⎟ ⎜ ⎜ RTtp ⎝ T R ⎝T ⎠ ⎠ +
ΔCP Ttp ln R T
(3)
In most chemical engineering applications, the first term of the right-hand side of eq 3 has the largest effect, and the next two terms with two opposite signs tend to compensate each other.27 Thus, ΔCP is negligible in comparison to the molar enthalpy of fusion and can be considered zero, thus ln x 2ideal =
ΔHfus ⎛ 1 1⎞ − ⎟ ⎜ R ⎝ Tm T⎠
(4)
then the ideal and experimental mole fractions are related by x2 = 1493
x 2ideal γ2
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where x2ideal is the saturation mole fraction of solute in an ideal solution, and x2 and γ2 obtained from experiments are used for the regression in the solubility prediction. The following equation yields the solubility S in terms of grams of solute per 100 grams of solvent using the saturation mole fraction, x2 x 2MW,solute
S=
The predictive accuracy was evaluated by the calculating root-mean-square error (RMSE) and coefficient of determination (R2) based on the experimental and the predicted solubility data. MATLAB software (MathWorks, Natick, MA) was used for all regression and related statistical analyses (at the 95 % confidence level).
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RESULTS AND DISCUSSION Physical Properties. Table 2 shows ΔHfus and Tm of both forms of L-phenylalanine based on the DSC and TGA experiments.29 The characteristic peaks in FTIR spectra of two forms are shown in Table 3 and Figure 1, in which the anhydrous and
× 100
(1 − x 2)MW,solvent
(6) −1
where MW is the molar mass, g·mol . The activity coefficient of the solute in the solvent, an indicator of the degree of nonideality, will be acquired using the UNIQUAC equation21,22 ϕi
ϕ θ z ln γi = ln + qi ln i + li − i xi xi 2 ϕi m
∑ xjlj
ΔHfus
j=1 m
− qi′ ln(∑ θ′jτji) + qi′ − qi′ ∑ j=1
Table 2. Dehydration Temperature, Td, Melting Point, Tm, and Heat of Fusion, ΔHfus, of L-Phe and L-Phe·H2O
k
j=1
form
θ′jτji (7)
L-phe·H2O
where xi are the mole fractions of component i, θi is the area fraction, and ϕi is the segment fraction that is similar to the volume fraction ϕi =
rx i i m
∑ j = 1 rjxj
θi =
qixi m
∑ j = 1 qjxj
θ′i =
Tm/K
J·g−1
361.02
549.75 552.15
111.78 126.62
L-phe
m
∑k = 1 θ′kτki
Td/K
Table 3. Assignment for Some Characteristic Vibrational Bands of L-Phe and L-Phe·H2O30 wavenumber/cm−1
qi′xi m
∑ j = 1 q′j xj
band assignment
L-phe
L-phe·H2O
benzene ring backbone vibrating
700 1074 1560 1625 747 2963 3031 1002 1224
701 1074 1559 1625 749 2960 3031
(8)
⎛ aij ⎞ τij = exp⎜ − ⎟ ⎝ T⎠
⎛ aji ⎞ τji = exp⎜ − ⎟ ⎝ T⎠
⎛z⎞ li = ⎜ ⎟(ri − qi) − (ri − 1) ⎝2⎠
CC stretching
(9) CN stretching
aij and aji are adjustable parameters of the UNIQUAC equation which can be obtained from the binary equilibrium data, z is assumed to be equal to 10, the dimensionless parameters r, q, and q′ are pure component constants, which depend on the molecular size of the solute and solvent molecules and can be calculated from van der Waals volume and area28
ri = qi =
hydrogen bond stretching
1197
Q vdwi 15.17
(10)
R vdwi 2.5 × 109
(11)
where Qvdwi and Rvdwi are the van der Waals molar volume and area of the molecule i. In the case that the van der Waals area and volume of the molecules are not available, the functional group approach by Fredenslund et al.24 can be used p
r=
∑ ni ·R i i=1
(12)
p
q=
∑ ni ·Q i
(13)
Figure 1. Characteristic absorption peaks of two solid forms of L-phenylalanine (above: monohydrate; below: anhydrate).
where p is the number of functional groups in the molecule and n is the repeating number of each functional group in the molecule. Ri and Qi are the dimensionless van der Waals area and volume of each functional group. q′ is equal to q, but for alcohols it differs, for example, r, q, and q′ for ethanol are 2.11, 1.97, and 0.92, respectively.8
monohydrate forms have characteristic absorption bands at wavenumbers ν of (1002 and 1197) cm−1, respectively.30 Optimization of Adjustable Parameters. Due to the lack of van der Waals area and volume of components, eqs 12 and 13 were used to calculate the dimensionless parameters r and q.
i=1
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Table 4. Volume and Area Parameters, Ri, Qi, and ni, of Functional Groups in L-Phe and L-Phe·H2O L-phe
Solubility Data. Based on the above information, the adjustable parameters are used to predict the solubility at different temperatures. The difference between the experimental and the predicted solubilities is an error defined by:
L-phe·H2O
functional group
Ri
Qi
ni
Ri
Qi
ni
CHNH2 (amine) COOH ACH (aromatic CH) ACCH2 (aromatic CH2) H2O
1.1417 1.3013 0.5313 1.0396
0.9240 1.2240 0.4000 0.6600
1 1 5 1
1.1417 1.3013 0.5313 1.0396 0.9200
0.9240 1.2240 0.4000 0.6600 1.4000
1 1 5 1 1
n ⎛ |S ⎞ exp − Scalc| ⎟⎟ · 100 %E = percent error = average ∑ ⎜⎜ S ⎝ ⎠ exp k=1
(15)
where Sexp and Scalc are the experimental and predicted solubility, respectively. The solubility prediction process is: (1) calculate the ideal mole fraction from eq 4 at a temperature of T, (2) value an initial mole fraction of the solute, (3) calculate the activity coefficient from UNIQUAC using the initial mole fraction, (4) calculate a new mole fraction for the solute and employ it as the new initial mole fraction of the solute using the calculated activity coefficient and the ideal mole fraction, and (5) repeat steps 3 and 4 until the mole fraction does not change anymore. The solubilities of L-phenylalanine anhydrous and monohydrate forms from measurements and predictions in different solvent systems at various temperatures are presented in Tables 7 to 14. The solubility of L-phenylalanine anhydrate and monohydrate agree quite well with the values given by Mohan et al. (Table 15);26 however, there exist large differences in the solubility of L-phenylalanine anhydrate in water + ethanol mixtures with different mass fractions of ethanol at 298.15 K between this work and the work of Nozaki and Tanford (Table 16).32 The solubility generally increases with the temperature (normal solubility), which suggests that heat is released when the crystals of L-phenylalanine anhydrate or monohydrate form.33 In addition, the solubility generally increases with a decrease in the content of acetone or ethanol in the solvent mixture, which confirms that acetone and ethanol can act as antisolvents for the drowning-out of L-phenylalanine from water. The influence of acetone and alcohol on the solubility of amino acids can be attributed to their effect on the hydrophobic interactions between solute and water. The promoted hydrophobicity of the solute with an increased amount of acetone and alcohol in the mixture will destabilize the intramolecular interactions between the solute and the water and will reduce the solubility accordingly.34
Table 5. Dimensionless Parameters, r, q, and q′, of Solvents solvent
r
q
q′
acetone ethanol water
2.5735 2.1100 0.9200
2.3360 1.9700 1.4000
2.3360 0.9200 1.0000
The values of n, R, and Q for L-phe and L-phe·H2O are given in Table 4. r and q of L-phe were then calculated as 6.1391 and 4.808, respectively, whereas those of L-phe·H2O were 7.0591 and 6.208, respectively. On the other hand, r, q, and q′ of the employed solvents were obtained from literature and presented in Table 5.31 The adjustable parameters (a12 and a21) for the mixture of solute and pure solvent were obtained using a few experimental solubility data and a nonlinear regression method. The optimization procedure was based on the minimization of the error between the calculated and the experimental values of activity coefficients d
min error =
parameters
∑ (γ2, k ,exp − γ2,k ,calc)
(14)
k=1
where d is the number of experimental data points and γ2,k,exp and γ2,k,calc are the experimental and calculated activity coefficients, respectively. Table 6 shows the adjustable parameters that were calculated from the solubility data. The adjustable parameters (a23 and a32) for the mixture of acetone and water were found to be 116.62 and −572.6, respectively, and those for ethanol and water were −567.3 and 63.810, respectively.
Table 6. Adjustable Parameters, a12 and a21, Obtained for Different Systems in the Range of Studied Concentrations, c L-phe
a
−1
L-phe·H2O
solvent
c/(g solute·(100 g solvent) )
a12a
acetone ethanol water
0.003−0.009 0.046−0.206 2.304−4.430
12.92 470.95 104.36
−1
a
a21
453.64 155.09 −129.60
c/(g solute·(100 g solvent) )
a12a
a21a
0.002−0.016 0.016−0.121 2.228−3.406
19.27 495.13 196.07
322.82 92.95 −105.30
Subscript 1 represents solute, and 2 represents solvent.
Table 7. Experimental Solubility, Sexp, of L-Phe in Water + Acetone Mixtures with Different Mass Fractions of Acetone, m1, in the Temperature Range from (283.15 to 323.15) K Sexp/(g solute·(100 g solvent)−1)
T K 283.15 293.15 298.15 303.15 313.15 323.15
m1 = 0 2.54 2.81 2.97 3.28 4.00 4.43
± ± ± ± ± ±
0.11 0.13 0.13 0.14 0.18 0.20
m1 = 0.08 2.37 2.67 2.85 3.07 3.88 4.33
± ± ± ± ± ±
0.11 0.12 0.13 0.14 0.17 0.19
m1 = 0.25 1.92 2.19 2.38 2.61 3.46 3.89
± ± ± ± ± ±
0.09 0.10 0.11 0.12 0.16 0.18
m1 = 0.45 1.20 1.40 1.60 1.74 2.57 2.71 1495
± ± ± ± ± ±
0.05 0.06 0.07 0.08 0.12 0.12
m1 = 0.65 0.43 0.47 0.56 0.68 1.18 1.22
± ± ± ± ± ±
0.02 0.02 0.03 0.03 0.05 0.05
m1 = 0.87 0.049 0.055 0.060 0.085 0.099 0.102
± ± ± ± ± ±
0.002 0.003 0.003 0.004 0.004 0.005
m1 = 1 0.0030 0.0043 0.0050 0.0056 0.0080 0.0089
± ± ± ± ± ±
0.0002 0.0002 0.0003 0.0003 0.0004 0.0005
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Table 8. Predicted Solubility, Scalc, of L-Phe in Water + Acetone Mixtures with Different Mass Fractions of Acetone, m1, in the Temperature Range from (283.15 to 323.15) K Scalc/(g solute·(100 g solvent)−1)
T K
m1 = 0
m1 = 0.08
m1 = 0.25
m1 = 0.45
m1 = 0.65
m1 = 0.87
m1 = 1
RMSE
R2
283.15 293.15 298.15 303.15 313.15 323.15
2.40 2.85 3.03 3.36 4.29 4.56
2.26 2.66 2.83 3.22 4.14 4.43
1.80 2.13 2.28 2.72 3.57 3.93
1.07 1.32 1.50 1.68 2.50 2.70
0.35 0.43 0.52 0.63 1.07 1.16
0.036 0.044 0.047 0.072 0.102 0.113
0.0036 0.0054 0.0064 0.0085 0.0113 0.0128
0.141 0.115 0.115 0.084 0.169 0.159
0.977 0.990 0.991 0.996 0.990 0.993
Table 9. Experimental Solubility, Sexp, of L-Phe in Water + Ethanol Mixtures with Different Mass Fractions of Ethanol, m2, in the Temperature Range from (283.15 to 323.15) K Sexp/(g solute·(100 g solvent)−1)
T K 283.15 293.15 298.15 303.15 313.15 323.15
m2 = 0 2.54 2.81 2.97 3.28 4.00 4.43
± ± ± ± ± ±
0.11 0.13 0.13 0.14 0.18 0.20
m2 = 0.08 1.95 2.31 2.49 2.64 3.13 3.91
± ± ± ± ± ±
m2 = 0.25
0.09 0.11 0.12 0.12 0.15 0.18
1.25 1.55 1.69 1.81 2.20 2.92
± ± ± ± ± ±
m2 = 0.45
0.06 0.07 0.08 0.09 0.10 0.14
0.74 0.96 1.07 1.20 1.46 1.96
± ± ± ± ± ±
0.04 0.05 0.05 0.05 0.07 0.09
m2 = 0.65 0.38 0.50 0.60 0.69 0.83 1.10
± ± ± ± ± ±
m2 = 0.87
0.02 0.03 0.03 0.04 0.04 0.05
± ± ± ± ± ±
0.091 0.127 0.153 0.195 0.283 0.367
0.005 0.006 0.007 0.009 0.013 0.017
m2 = 1 0.046 0.067 0.074 0.106 0.150 0.179
± ± ± ± ± ±
0.003 0.004 0.004 0.005 0.007 0.009
Table 10. Predicted Solubility, Scalc, of L-Phe in Water + Ethanol Mixtures with Different Mass Fractions of Ethanol, m2, in the Temperature Range from (283.15 to 323.15) K Scalc/(g solute·(100 g solvent)−1)
T K
m2 = 0
m2 = 0.08
m2 = 0.25
m2 = 0.45
m2 = 0.65
m2 = 0.87
m2 = 1
RMSE
R2
283.15 293.15 298.15 303.15 313.15 323.15
2.40 2.85 3.03 3.36 4.29 4.56
1.92 2.34 2.49 2.78 3.15 3.83
1.29 1.60 1.72 1.94 2.15 2.78
0.79 1.01 1.12 1.27 1.48 1.94
0.43 0.54 0.63 0.73 0.90 1.18
0.11 0.15 0.17 0.21 0.30 0.38
0.025 0.043 0.047 0.052 0.065 0.112
0.113 0.097 0.101 0.167 0.087 0.203
0.981 0.990 0.990 0.979 0.995 0.983
Table 11. Experimental Solubility, Sexp, of L-Phe·H2O in Water + Acetone Mixtures with Different Mass Fractions of Acetone, m1, in the Temperature Range from (278.15 to 308.15) K Sexp/(g solute·(100 g solvent)−1)
T K 278.15 283.15 293.15 298.15 303.15 308.15
m1 = 0 2.23 2.30 2.71 2.86 3.21 3.41
± ± ± ± ± ±
m1 = 0.08
0.10 0.11 0.12 0.13 0.14 0.15
1.86 2.05 2.50 2.70 3.05 3.28
± ± ± ± ± ±
m1 = 0.25
0.09 0.09 0.11 0.12 0.13 0.14
1.24 1.47 1.91 2.11 2.50 2.86
± ± ± ± ± ±
m1 = 0.45
0.06 0.07 0.09 0.10 0.11 0.12
0.65 0.79 1.01 1.10 1.46 1.91
± ± ± ± ± ±
0.03 0.04 0.05 0.05 0.06 0.08
m1 = 0.65 0.155 0.220 0.344 0.396 0.607 0.739
± ± ± ± ± ±
0.008 0.009 0.015 0.017 0.027 0.033
m1 = 0.87 0.0040 0.0065 0.0090 0.0250 0.0500 0.0820
± ± ± ± ± ±
0.0002 0.0003 0.0004 0.0011 0.0022 0.0037
m1 = 1 0.0020 0.0025 0.0037 0.0043 0.0052 0.0061
± ± ± ± ± ±
0.0001 0.0001 0.0002 0.0002 0.0003 0.0003
Table 12. Predicted Solubility, Scalc, of L-Phe·H2O in Water + Acetone Mixtures with Different Mass Fractions of Acetone, m1, in the Temperature Range from (278.15 to 308.15) K Scalc/(g solute·(100 g solvent)−1)
T K
m1 = 0
m1 = 0.08
m1 = 0.25
m1 = 0.45
m1 = 0.65
m1 = 0.87
m1 = 1
RMSE
R2
278.15 283.15 293.15 298.15 303.15 308.15
2.17 2.35 2.76 2.86 3.27 3.47
1.91 2.08 2.51 2.63 3.09 3.28
1.28 1.52 1.82 2.00 2.42 2.70
0.59 0.74 0.93 1.02 1.35 1.70
0.12 0.18 0.28 0.35 0.52 0.63
0.002 0.003 0.013 0.038 0.046 0.091
0.0019 0.0022 0.0031 0.0037 0.0044 0.0050
0.069 0.059 0.094 0.098 0.189 0.260
0.994 0.996 0.993 0.993 0.979 0.965
content of organic solvent, with the average root-mean-square error (RMSE) and coefficient of determination (R2) as 0.1187 and 0.9883, respectively. However, in the case of mass fractions
As shown in Tables 7 to 14, the UNIQUAC model can work well for the prediction of the solubility of the pseudopolymorphic system studied in water and its binary mixtures with a low 1496
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Table 13. Experimental Solubility, Sexp, of L-Phe·H2O in Water + Ethanol Mixtures with Different Mass Fractions of Ethanol, m2, in the Temperature Range from (278.15 to 308.15) K Sexp/(g solute·(100 g solvent)−1)
T m2 = 0
K 278.15 283.15 293.15 298.15 303.15 308.15
2.23 2.30 2.71 2.86 3.21 3.41
± ± ± ± ± ±
m2 = 0.08
0.10 0.11 0.12 0.13 0.14 0.15
1.51 1.67 2.07 2.32 2.54 2.80
± ± ± ± ± ±
m2 = 0.25
0.07 0.08 0.09 0.10 0.11 0.12
0.84 1.01 1.39 1.57 1.77 2.04
± ± ± ± ± ±
m2 = 0.45
0.04 0.05 0.06 0.07 0.08 0.09
0.45 0.57 0.85 0.99 1.17 1.36
± ± ± ± ± ±
0.02 0.03 0.04 0.04 0.05 0.06
m2 = 0.65 0.200 0.283 0.415 0.491 0.635 0.754
± ± ± ± ± ±
m2 = 0.87
0.009 0.010 0.017 0.020 0.027 0.031
0.041 0.057 0.097 0.135 0.176 0.215
± ± ± ± ± ±
0.002 0.003 0.004 0.006 0.008 0.009
m2 = 1 0.0160 0.0290 0.0590 0.0720 0.0920 0.1210
± ± ± ± ± ±
0.0009 0.0013 0.0025 0.0030 0.0040 0.0050
Table 14. Predicted Solubility, Scalc, of L-Phe·H2O in Water + Ethanol Mixtures with Different Mass Fractions of Ethanol, m2, in the Temperature Range from (278.15 to 308.15) K Scalc/(g solute·(100 g solvent)−1)
T K
m2 = 0
m2 = 0.08
m2 = 0.25
m2 = 0.45
m2 = 0.65
m2 = 0.87
m2 = 1
RMSE
R2
278.15 283.15 293.15 298.15 303.15 308.15
2.17 2.35 2.76 2.86 3.27 3.47
1.44 1.65 2.01 2.24 2.52 2.73
0.87 1.03 1.31 1.48 1.68 1.94
0.51 0.63 0.88 1.00 1.14 1.30
0.24 0.35 0.47 0.55 0.61 0.77
0.05 0.08 0.12 0.14 0.19 0.21
0.002 0.003 0.040 0.044 0.052 0.066
0.089 0.061 0.099 0.098 0.076 0.106
0.985 0.994 0.988 0.990 0.995 0.992
Table 15. Comparing Experimental Solubility, Sexp, of L-Phe and L-Phe·H2O in Water at Different Temperatures with Those from Reference 26 Sexp,anhydrate/ (g solute·(100 g solvent)−1)
T K
this work
278.15 283.15 293.15 298.15 303.15 308.15 313.15 323.15
2.54 2.81 2.97 3.28 4.00 4.43
ref 26
2.71 2.92 3.12 3.36 3.69 4.29
S exp,monohydrate/ (g solute·(100 g solvent)−1) this work
ref 26
2.23 2.30 2.71 2.86 3.21 3.41
2.08 2.53 2.71 2.97 3.29 3.78 4.69
Figure 2. Experimental solubility of two forms in water at various temperatures (○, anhydrate; □, monohydrate).
Table 16. Comparing Experimental Solubility, Sexp, of L-Phe in Water + Ethanol Mixtures with Different Mass Fractions of Ethanol, m2, at 298.15 K with Those from Reference 32
them is about 308.9 K, which is almost accordance with the results of Mohan et al.26 The anhydrous form is the thermodynamically favored form (stable form) above the transition temperature, whereas the monohydrate form is thermodynamically favored form (stable form) below the transition temperature.
Sexp,anhydrate/(g solute·(100 g solvent)−1) m2
this work
ref 32
0 0.08 0.25 0.45 0.65 0.87 1
2.97 2.49 1.69 1.07 0.60 0.153 0.074
2.79 2.33 1.75 1.50 1.35 0.829 0.153
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CONCLUSIONS Solubility information can be used to distinguish the difference between different solid forms of an API. This contribution has presented the experimental solubility of L-phenylalanine anhydrous and monohydrate forms in various solvents at different temperatures. The solubility is remarkably dependent upon the temperature and the content of acetone or ethanol in the solution. Meanwhile, the semiempirical UNIQUAC model with adjustable parameters that can be estimated from the binary equilibrium data has been demonstrated to predict the solubility of two forms in water and its binary mixtures with low content of organic solvent well. All of these results shall afford a better understanding and crystallization process control of this enantiotropic pseudopolymorphic system.
of ethanol and acetone higher than 0.65, the deviation between experimental and calculated data ranges from 37 % at 323.15 K to 46 % at 298.15 K, indicating that the model seemed unavailable. Transition Temperature. The transition point between two forms can thus be determined by plotting their solubility in pure water against temperature, as shown in Figure 2. The solubility curves expose an enantiotropic nature of the two forms of L-phenylalanine, and the transition point between 1497
dx.doi.org/10.1021/je201354k | J. Chem. Eng. Data 2012, 57, 1492−1498
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Article
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AUTHOR INFORMATION
Corresponding Author
*Fax: 86-510-85917763. E-mail:
[email protected]. Funding
This work was supported by the National Natural Science Foundation of China (Nos. 21176102 and 21176215), the Scientific Research Foundation for Returned Chinese Scholars, and the Fundamental Research Funds for the Central Universities (No. JUSRP30904). Notes
The authors declare no competing financial interest.
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