Measurement and Correlation of the Solubility of l-Carnitine in

May 14, 2014 - ABSTRACT: A gravimetric method was used to determine the solubility of L-carnitine in different pure solvents and ethanol− acetone so...
0 downloads 0 Views 2MB Size
Download PDF
Article pubs.acs.org/jced

Measurement and Correlation of the Solubility of L‑Carnitine in Different Pure Solvents and Ethanol−Acetone Solvent Mixture Dengqiong Sun,†,‡ Riju Ren,†,‡ Weiqiang Dun,†,‡ Haihong Zhang,§ Lijun Zhao,§ Li Zhang,§ Wenqing Zhang,§ and Junbo Gong*,†,‡ †

School of Chemical Engineering and Technology, State Key Laboratory of Chemical Engineering, Tianjin University, Tianjin, 300072 People’s Republic of China ‡ The Co-Innovation Center of Chemistry and Chemical Engineering of Tianjin, Tianjin University Tianjin, 300072 People’s Republic of China § Northeast Pharmaceutical Group Co., Ltd., Shenyang, 110026 China S Supporting Information *

ABSTRACT: A gravimetric method was used to determine the solubility of L-carnitine in different pure solvents and ethanol− acetone solvent mixture. The solubility in different pure solvents were then correlated by the modified Apelblat equation, λh equation and the modified van’t Hoff equation, with the modified van’t Hoff equation presenting the best consistence. Meanwhile, to illustrate the effect of ethanol or acetone on the change of the solubility, a new parameter defined as influence coefficient was introduced and the coefficient of acetone depending on temperature and molar faction of acetone was depicted. In addition, the changes for enthalpy, entropy and Gibbs free energy were calculated by the modified van’t Hoff equation. It can be drawn that the changes for enthalpy and entropy in a solvent mixture decrease to a minimum before consequent increasing with increasing molar faction of acetone. Furthermore, the change for Gibbs free energy shows a linear relationship with natural logarithm of the solubility.

1. INTRODUCTION L-Carnitine (β-hydroxyl-γ-trimethylaminobutyrate; Figure 1), a small highly polar zwitterionic molecule, plays an important

sport beverages. Besides, it is used in treating chronic renal failure and relevant deficiency.3 In industrial manufacturing, for the poor efficiency of the extraction from animal viscera, L-carnitine is mainly produced through asymmetric synthesis4 followed by solution crystallization and further recrystallization which are the key steps to obtain the desired products. The solubility is an important parameter to be determined for a series of processes depending on its recognition, especially in the event of crystallization.5 Therefore, the measurement of the solubility of L-carnitine in different solvent systems is necessary. However, the solubility of L-carnitine has rarely been reported in the literature so far. In this work, the solubility data of L-carnitine in methanol, ethanol, isopropanol, acetone, and an ethanol−acetone mixture were determined by the gravimetric method within the

Figure 1. Molecular structure of L-carnitine.

role in the human metabolism via translocation of long-chain fatty acids across the mitochondrial inner membrane.1,2 With this special biochemical property, it is widely used in pharmaceutical and food industries, for instance, as a weightloss drug against the obesity caused by disorder of lipid metabolism and as functional food additive in milk powder and © 2014 American Chemical Society

Received: February 9, 2014 Accepted: May 6, 2014 Published: May 14, 2014 1984

dx.doi.org/10.1021/je500078n | J. Chem. Eng. Data 2014, 59, 1984−1990

Journal of Chemical & Engineering Data

Article

where m1, m2, and m3 represent the mass of L-carnitine and each solvent used in the solution, respectively, and M1, M2, and M3 are the molar mass of L-carnitine and each solvent.

temperature range from 278.15 K to 348.15 K. The measured solubility in different pure solvents was then correlated by the modified Apelblat equation, the λh equation, and the modified van’t Hoff equation. Meanwhile, the solubility data in the solvent mixture was correlated by the modified van’t Hoff equation and the Jouyban−Acree model. To illustrate the effect of ethanol and acetone on the change of the solubility, a new parameter defined as the influence coefficient was introduced. Also, the changes for the enthalpy, entropy and Gibbs free energy were estimated based on the solubility data.

3. RESULTS AND DISCUSSION 3.1. Solubility of L-Carnitine in Pure Solvents. The experimental molar fraction solubility of L-carnitine in methanol, ethanol, isopropanol, and acetone at the temperature range from 278.15 K to 348.15 K were presented in Table 1 along with the uncertainty. In this work, the solubility data of L-carnitine in different solvents were correlated by several models so as to quantitatively describe the solid−liquid equilibrium. In order to correlate the solubility with temperature, the modified Apelblat equation and semiempirical Buchowski−Ksiazczak λh equation were used. The modified Apelblat equation6 is written as B ln x1 = A + + C ln(T /K) (3) T /K

2. EXPERIMENTAL SECTION 2.1. Materials. L-Carnitine was supplied by Northeast Pharmaceutical Group Co., Ltd. (Shenyang, China) with the purity of the mass fraction higher than 0.998. Organic solvents including methanol, ethanol, isopropanol and acetone were acquired from Tianjin Kewei Reagent Co. (Tianjin, China) and the mass fractions of purity were higher than 0.995. All chemicals were used without further purification, and the physical properties are listed in the Supporting Information. 2.2. Apparatus and Methods. Before the measurements, L-carnitine was dried in a vacuum oven at 323.15 K to avoid the moisture uptake. Water content was then obtained lower than 0.1 % by a volumetric KF titrator (type: V20, Mettler Toledo, Switzerland). In the measurement of the solubility of L-carnitine, a gravimetric method was adopted. The slurry of L-carnitine was prepared by adding excess solute into each solvent in an Erlenmeyer flask with a working volume of 50 mL. The slurry was then allowed to equilibrate for more than 8 h under a desired temperature and agitation. After then, the agitator was turned off and the suspension was allowed to settle for more than 4 h to ensure solid phase precipitated to the bottom before sampling. A portion of the supernatant was filtered with an organic membrane (0.22 μm) and quickly transferred to three 5 mL beakers of given masses. The beakers were then weighed and subsequently placed into a vacuum oven at the temperature 323.15 K until constant dry weight. The saturation process had already been studied before the measurement. Four pure solvents and two different proportions of solvent mixtures with exceed L-carnitine added were kept at 293.15 K for 2 days and the compositions were measured every 1 h. It turned out to be that the dissolution process quickly reached equilibrium in about 2 h, and nearly a constant composition after 2 h. So we can confirm that by using the gravimetric method, the saturation can be reached in the measurement of the solubility of L-carnitine. The temperature and agitation were respectively determined by the built-in air bath and oscillator of a thermostatic oscillator (type HNY-200R, Tianjin Ounuo Instrument Co. Ltd., China). All of the masses in the experiment were determined by an electronic balance (type ML204/02, Mettler Toledo, Switzerland) with an accuracy of ± 0.0001 g. The molar fraction solubility of L-carnitine (x1) in pure solvents and solvent mixture can be calculated respectively by the following equations: m1/M1 x1 = m1/M1 + m2 /M 2 x1 =

m1/M1 m1/M1 + m2 /M 2 + m3 /M3

where x1 is the molar fraction solubility of L-carnitine; T represents the experimental temperature; and A, B, and C are the model parameters determined by the experimental solubility data, listed in the Supporting Information. The λh equation, which was originally developed by Buchowski et al.,7 correlates the molar fraction of the solubility with temperature by using only two parameters. The equation is defined as follows: ⎛1 ⎛ λ(1 − x1) ⎞ 1 ⎞ ln⎜1 + ⎟ ⎟ = λh⎜ − x1 Tm ⎠ ⎝T ⎠ ⎝

(4)

where x1 is the molar fraction solubility of L-carnitine, T and Tm are equilibrium temperature and melting temperature of the solute, respectively, and λ and h are the model parameters, which were presented in the Supporting Information. For real solutions, the standard van’t Hoff equation8 describes the relationship between the molar fraction solubility of a solute and temperature by taking the solvent effect into account. The relationship is defined as follows:

ln x1 =

ΔSd ΔHd − R RT

(5)

where x1 is the molar fraction solubility of L-carnitine and R and T represent the gas constant and the equilibrium temperature (T/K). ΔHd and ΔSd are respectively the standard changes for enthalpy and entropy of the dissolving process. The measurement and experimental errors would propagate into a consistent pattern on enthalpy−entropy plots. The slope of a line in which these errors are distributed has been called error slope. In order to describe the error slope, a so-called harmonic temperature Thm is introduced.The parameter was first used by RR Krug et.al.,9,10 defined as follows: n Thm = n ∑1 (1/T ) (6) For any n ≥ 2 experimental temperatures, the error slope is equal to Thm when enthalpies and entropies are estimated by ordinary linear regression.9,10 The modified van’t Hoff equation can be defined as11

(1)

−ΔHd ∂ln x1 = R ∂(1/T − 1/Thm)P

(2) 1985

(7)

dx.doi.org/10.1021/je500078n | J. Chem. Eng. Data 2014, 59, 1984−1990

Journal of Chemical & Engineering Data

Article

Equation 7 can be written as follows:

Table 1. Values of Solubility of L-Carnitine in Different Pure Solvents with Temperatures Ranging from 278.15 K to 348.15 K under Atmospheric Pressurea

ln x1 = A + B(1/T − 1/Thm)

in which A and B are the model parameters, which are listed in the Supporting Information. The relative deviation (RD), the relative average deviation (RAD) and the root-mean-square deviation (RMSD) were used to assess the accuracy and predictability of correlation models. The equations are defined as follows:

102RD solvent methanol

ethanol

isopropanol

acetone

T/K 278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15 333.15 278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15 333.15 338.15 343.15 348.15 278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15 333.15 278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15 278.15 283.15

103x1 96.244 98.234 103.688 110.635 118.964 127.123 135.565 144.165 154.137 165.167 177.535 189.346 39.164 40.825 43.106 45.224 48.330 51.747 54.583 57.644 60.725 64.804 69.750 74.721 80.095 85.518 91.938 1.463 1.810 2.233 2.708 3.404 4.173 5.266 6.295 7.779 9.831 11.835 14.427 0.071 0.081 0.091 0.106 0.117 0.131 0.143 0.157 0.175 0.195 0.220 0.071 0.081

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

1.652 0.974 1.536 1.963 1..721 1.534 0.995 1.741 1.236 2.031 1.275 1.362 0.335 0.279 0.578 0.623 0.918 0.295 0.837 0.623 0.798 1.427 0.785 0.912 0.845 1.297 1.164 0.037 0.042 0.025 0.065 0.078 0.056 0.039 0.047 0.096 0.062 0.136 0.095 0.001 0.003 0.002 0.005 0.002 0.001 0.002 0.004 0.003 0.001 0.006 0.001 0.003

eq 3

eq 8

1.138 −1.173 −0.958 −0.291 0.688 0.739 0.408 −0.285 −0.446 −0.33 0.193 0.063 −0.81 −0.44 0.394 −0.051 0.868 1.633 0.702 −0.238 −1.488 −1.417 −0.423 0.185 0.969 1.491 2.827 −26.146 3.259 12.831 10.916 8.005 2.083 0.127 −5.461 −5.54 −1.314 −0.364 3.283 −2.995 −0.942 0.163 3.05 1.345 1.273 −1.29 −2.547 −1.814 −0.768 1.958 −2.995 −0.942

−1.024 0.645 −0.905 −0.842 0.545 1.315 −0.821 1.214 −0.721 −0.824 0.546 −0.842 −0.412 0.255 −0.591 −0.675 0.725 −0.636 0.544 −0.258 −1.255 0.535 −0.788 0.076 0.557 −0.425 0.358 0.652 −1.275 0.684 1.157 −0.857 0.679 1.258 −1.388 0.875 1.668 −1.257 1.357 −0.436 −0.459 0.243 −2.464 0.912 −1.146 −1.1 −2.112 1.268 −0.265 2.258 −0.436 −0.459

(8)

RD =

x1 − x1cal x1 1

RAD =

∑ N

(x1 − x1cal)/N x1

⎡1 RMSD = ⎢ ⎢⎣ N

⎤1/2 cal 2 ⎥ ( x − x ) ∑ 1 1 ⎥⎦ N

(9)

(10)

1

(11)

where N is the number of experimental points obtained in each solvent which equals the number of temperatures used, xcal 1 represents the values of solubility calculated, and x1 is the solubility measured. Equations 9, 10, and 11 were defined to calculate the RD, RAD, and RMSD, respectively. It can be concluded from the Supporting Information that the solubility data of L-carnitine in four different solvents are well correlated by the modified Apelblat equation, the λh equation, and the modified van’t Hoff equation. For all models used in this work, the modified van’t Hoff equation shows the best correlation in describing the dependence of solubility on temperature with the RAD of methanol, ethanol, isopropanol, and acetone: 0.85 %, 0.54 %, 1.09 %, and 1.15 %, respectively. All values of RAD are lower than 1.2 %. To illustrate in detail, Figures 2 and 3 were presented, from which we can see that the solubility of L-carnitine in these selected pure organic solvents is temperature-dependent, increasing with increasing temperature. It can also be drawn that the solubility of L-carnitine in highly polar protic solvents is higher than that in weakly polar aprotic solvents, which is corresponding to its highly polar zwitterionic character.

a The solubility data is correlated by the modified Apelblat equation (eq 3) and the modified van’t Hoff equation (eq 8). Standard uncertainty ur (T) = 0.1 K, ur (P) = 0.05, ur (x) = 0.016.

Figure 2. Modified Apelblat plots of x1 versus T in different pure organic solvents: (■) methanol, (●) ethanol, (▲) isopropanol, and (▼) acetone. 1986

dx.doi.org/10.1021/je500078n | J. Chem. Eng. Data 2014, 59, 1984−1990

Journal of Chemical & Engineering Data

Article

Figure 4. Solubility of L-carnitine (x1) depending on temperature T and the molar fraction of acetone (XA) in acetone + ethanol solvent mixture: (○) XA = 0.0817, (Δ) XA = 0.1655, (▽) XA = 0.2530, (□) XA = 0.3463, (◇) XA = 0.4431, (■) XA = 0.5432, (●) XA = 0.6479, (▲) XA = 0.7590, and (▼) XA = 0.8761.

Figure 3. Modified van’t Hoff plots of ln x1 versus 1/T − 1/Thm in different pure organic solvents: (□) methanol, (○) ethanol, (Δ) isopropanol, and (▽) acetone.

3.2. Solubility of L-Carnitine in Ethanol−Acetone Solvent Mixture. A solvent mixture is widely used in the crystallization process of L-carnitine in industry, so it is of great importance to measure the solubility of L-carnitine in solvent mixture. According to the analysis of pure solvents, we choose ethanol and acetone as the solvent mixture, because they are respectively highly polar protic and weakly polar aprotic solvents. In this work, the solubility data of L-carnitine in solvent mixture consist of different proportions of ethanol− acetone was measured at different temperatures, which are listed in Table 2 along with uncertainty. The content is shown in Figure 4 by correlating the solubility data to the equilibrium temperature (T) and the molar fraction of acetone (XA). From Table 2 and Figure 4, it can be concluded that the solubility data in ethanol−acetone solvent mixture depends on both temperature and XA. The solubility of L-carnitine decreases with increasing XA for a fixed value of temperature, while for a fixed value of XA, the solubility increases with increasing temperature. In this work, the solubility data in solvent mixture was correlated by the Jouyban−Acree model,12 which was defined as eq 12.

ln x1 = w2 ln x E + w3 ln xA +

w2w3 i = 0 ·∑ Ji (w2 − w3)i T 2

(12)

where x1, xE, and xA respectively are the molar fraction solubility data of L-carnitine in ethanol−acetone solvent mixture, pure ethanol, and acetone and w2 and w3 are the mass fraction of ethanol and acetone in the absence of Lcarnitine. The Ji terms are the model parameters determined by eq 13: [ln x1 − (w2 ln x E + w3 ln xA)]T w2w3 = J0 + J1(w2 − w3) + J2 (w2 − w3)2

(13)

The model parameters J0, J1, and J2 were determined by correlating the value of [ln x1 − (w2 ln xE + w3 ln xA)]T/w2w3 to (w2 − w3)i, i = 0, 1, 2. The values of J0, J1, and J2 were presented in the Supporting Information. At constant temperature, the solubility of a solute is determined by the chemical properties of the solvent as well as the solute−solute, solute−solvent, and solvent−solvent interactions. In solvent mixture, these interactions are

Table 2. Solubility Data of L-Carnitine in Different Proportions of Ethanol−Acetone Solvent Mixture at Different Temperaturesa) 103x1

a

XA

T

0.0817

0.1655

0.2530

0.3463

0.4431

0.5432

0.6479

0.7597

0.8761

278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15

35.195 35.756 37.832 40.375 42.068 45.360 47.331 50.187 52.090 55.725 61.791

31.248 32.132 33.991 35.577 37.115 38.419 40.556 42.861 44.908 46.379 48.805

24.733 25.323 26.709 27.433 28.791 29.486 31.057 32.010 34.044 36.086 38.409

17.707 18.450 19.442 19.405 19.953 20.480 21.353 22.412 23.494 25.950 28.620

11.929 12.259 12.353 12.863 13.147 13.694 14.023 14.617 15.935 17.595 20.183

7.320 7.382 7.651 7.582 7.692 8.105 8.391 8.738 8.987 9.618 10.205

4.103 4.229 4.307 4.496 4.624 4.707 4.767 4.890 5.103 5.253 5.391

1.354 1.447 1.537 1.649 1.777 1.930 2.085 2.230 2.375 2.530 2.746

0.504 0.546 0.587 0.639 0.693 0.757 0.823 0.885 0.951 1.020 1.115

XA is the molar fraction of acetone in the absence of L-carnitine. Standard uncertainty ur (T) = 0.1 K, ur (P) = 0.05, ur (x) = 0.016. 1987

dx.doi.org/10.1021/je500078n | J. Chem. Eng. Data 2014, 59, 1984−1990

Journal of Chemical & Engineering Data

Article

related to the consumption and release of heat (ΔHd), which is depending on the nature of interactions of solvent−solvent, solvent−solute and solute−solute. Entropy change (ΔSd) is an indicator of the disorder. In this work, the modified van’t Hoff equation was used to calculate ΔHd of the dissolving process, obtained by the slope (−ΔHd/R) in the linear relation of lnx1 versus (1/T − 1/Thm). It is known that the solubility (ln x1) is related to the standard molar Gibbs energy of solvation. In an ideal solvation, the standard Gibbs energy in a solvent mixture depends on the molar fraction of each solvent and the solubility in the solvent mixture. So as to determine the standard Gibbs energy change (ΔGd), the parameter γ was introduced to describe it:9,10

significant that must be taken into consideration. In order to analyze the contribution of ethanol and acetone to the change of solubility in solvent mixture, a new parameter δ was defined as follows: δE =

ln x1 − (XE ln x E + XA ln xA ) = −δA ln x E − ln xA

(14)

in which δE and δA were respectively used to calculate the influence coefficient of ethanol and acetone of the solvent mixture in the local region. x1, xE, and xA denote the solubility data of L-carnitine in the ethanol−acetone solvent mixture, pure ethanol, and acetone, respectively, and XE and XA are the molar fraction of ethanol and acetone in the absence of L-carnitine. Figure 5 presents the influence coefficient of acetone (δA) which is determined by the temperature (T/K) and the molar

ΔGd = −RThm*γ

(16)

in which, γ represents the intercept of the linear relation of ln x1 versus (1/T − 1/Thm). It is easily known that the enthalpy can be obtained by the traditional method, while the standard Gibbs energy change is redefined. This is because ΔGd does not depend on the value obtained at the specific temperature only, whereas, it is determined by all the solubility data. The entropy change for the dissolution process is then obtained at Thm: ΔSd =

ΔHd − ΔGd Thm

(17)

The relative contributions of the enthalpy (rH) and entropy (rS) to the standard Gibbs free energy of the dissolving process were respectively defined as follows:

Figure 5. Influence coefficient of acetone (δA) depending on the temperature T and the molar fraction of acetone (XA) in the ethanol− acetone solvent mixture.

|ΔHd| |ΔHd| + |ThmΔSd|

(18)

rS =

|ThmΔSd| |ΔHd| + |ThmΔSd|

(19)

The solubility data of L-carnitine in solvent mixture was correlated by the modified van’t Hoff equation. Figure 6 shows the modified van’t Hoff plots of ln x1 versus (1/T − 1/Thm) in

fraction of acetone (XA) in the ethanol−acetone solvent mixture. It can be obviously concluded that the acetone performs as an antisolvent, for the fact that the coefficient δA is always less than zero. The values of δA slightly decrease with increasing temperature and a maximum δA would be obtained with increasing XA, where the influence of acetone on the solubility of L-carnitine in the solvent mixture reaches the maximum. The maximum means a balance between the concentration of acetone and solute molecules. Exceed acetone cannot reduce the solubility as expected, thus decreasing the influence coefficient δA. The combined nearly ideal binary solvent/Redlich (CNIBS/ R-K) equation13,14 can be used to describe the relevance of δ and the molar fraction of ethanol (XE), which is shown as follows: δ E = A 0 + A1 XE + A 2 XE 2 + A3XE 3 + A4 XE 4

rH =

(15)

where A0, A1, A2, A3, and A4 are parameters of this model and the values of them are listed in the Supporting Information. The nonlinear variation of δ with XE was determined by the preferential solvation of the solute in the solvent mixture.15 3.3. Thermodynamic Analysis in the Ethanol−Acetone Solvent Mixture. The solvation process is closely

Figure 6. Modified van’t Hoff plots of ln x1 versus (1/T − 1/Thm) in acetone (XA) + ethanol solvent mixture: (○) XA = 0.0817, (Δ) XA = 0.1655, (▽) XA = 0.2530, (□) XA = 0.3463, (◇) XA = 0.4431, (×) XA = 0.5432, (☆) XA = 0.6479, (●) XA = 0.7590, (■) XA = 0.8761. 1988

dx.doi.org/10.1021/je500078n | J. Chem. Eng. Data 2014, 59, 1984−1990

Journal of Chemical & Engineering Data

Article

ethanol−acetone solvent mixture. The thermodynamic parameters of L-carnitine in dissolving process were calculated in different pure solvents and solvent mixture with different proportions. The results are presented in Table 3. It can be Table 3. Thermodynamic Parameters of Dissolving Process (at Thm = 302.32K) ΔHd solvent methanol isopropanol ethanol + acetone

XA

0.0000 0.0817 0.1655 0.2530 0.3463 0.4431 0.5432 0.6479 0.7597 0.8761 1.0000

−1

ΔGd −1

kJ·mol

kJ·mol

9.85 32.22 9.85 8.39 6.89 6.50 6.40 7.01 4.87 4.06 10.82 11.45 16.73

5.10 13.50 7.14 7.79 8.16 8.81 9.67 10.69 12.05 13.47 15.72 17.06 22.55

ΔSd J·mol−1·K−1

rH %

15.71 61.94 8.96 1.99 −4.22 −7.64 −10.82 −12.16 −23.73 −31.15 −16.23 −18.55 −19.22

67.47 63.25 78.42 93.30 84.36 73.76 66.17 65.60 40.45 30.10 68.80 67.12 74.22

Figure 7. Linear relationship of ΔGd (kJ mol−1) versus the natural logarithm of molar faction solubility ln x1 at Thm = 302.32 K.

selection of solvent and antisolvent when dealing with other substances. The analysis of solubility in ethanol−acetone solvent mixture manifests a close relationship between the dissolving process and solvent−solvent interactions. First, the Jouyban−Acree model was very consistent with the solubility data. Then, a parameter δ was defined to calculate the influence coefficient of acetone in solvent mixture in the local region, and the result is that the value of δA decreases with increasing temperature and a maximum δA would be obtained with increasing XA. The maximum is corresponding to the max influence of acetone on the solubility of L-carnitine in ethanol−acetone solvent mixture. Finally, the values of δA were well correlated by the combined nearly ideal binary solvent/Redlich (CNIBS/R-K) equation. In the part of thermodynamic analysis, it is indicated that the changes for the enthalpy and the entropy of the dissolving process in the solvent mixture consist of different proportions decrease to a minimum before consequent increasing with increasing XA, while the values of ΔGD are found to have a linear relationship with the natural logarithm of the solubility ln x1 . Overall, the experimental solubility and parameters in this study will be used for the crystallization of L-carnitine in industry and conclusions obtained from this work will make a contribution to the analysis of some other substances.

concluded that dissolving process of L-carnitine in the solvent mixture are endothermic and a minimum of the enthalpy will be obtained with increasing molar fraction of acetone (XA). The change for Gibbs free energy shows variation of increasing with increasing XA which is corresponding to increasing solubility. In ethanol−acetone solvent mixture, the entropy change decreases to a minimum before consequent increasing with increasing XA. The positive entropy change indicates that solubilization process is entropy unfavorable,16 while the negative entropy change is due to the increased order in solutions.17 The functional groups present in L-carnitine and the solvent may be the key factors to determine the entropy change. Because of the existence of the groups like −N−, −OH, and −COO−−, Lcarnitine may involve different forces such as hydrogen bond, hydrophobic interaction and electrostatic force in the dissolving process.18 Considering the relative contribution of enthalpy change (rH), a minimum can be obtained with increasing XA, where the dissolving process is least affected by the dissolution enthalpy, and most closely related to the entropic term. ΔGD of L-carnitine at Thm = 302.32K in different solvent systems was plotted with the natural logarithm of the molar fraction solubility (ln x1) in Figure 7. It is indicated that a linear relationship exists between ΔGD and ln x1, which shows that lower values of ΔGD corresponds to a more favorable process of dissolution and to higher solubility.19

■

ASSOCIATED CONTENT

S Supporting Information *

Chemical properties (mass purity, molar mass, and molar volume) of pure materials, experimental solubility, calculated parameters for different models, and predicted values of the mixing properties in different pure solvents and solvent mixture consist of different proportions. This material is available free of charge via the Internet at http://pubs.acs.org.

4. CONCLUSION A gravimetric method was used to determine the solubility data of L-carnitine in methanol, ethanol, isopropanol, and acetone pure solvents and ethanol−acetone mixed solvents at the temperature ranging from 278.15 K to 348.15 K. For different pure solvents, the solubility data was well correlated by the modified Apelblat equation, λh equation, and the modified van’t Hoff equation. The modified van’t Hoff equation presents the best consistency in general. It can be concluded that the solubility of L-carnitine in protic solvents is better than in aprotic solvents and L-carnitine is less soluble in the solvent with strong polarity. This knowledge will contribute to the

■

AUTHOR INFORMATION

Corresponding Author

*Tel.: 86-22-27405754. Fax: +86-22-27374971. E-mail: junbo_ [email protected]. Funding

We are grateful for the financial support of the National Natural Science Foundation of China (No. NNSFC 21176173), and 1989

dx.doi.org/10.1021/je500078n | J. Chem. Eng. Data 2014, 59, 1984−1990

Journal of Chemical & Engineering Data

Article

(11) Pacheco, D. P.; Manrique, Y. J.; Martínez, F. Thermodynamic study of the solubility of ibuprofen and naproxen in some ethanol+ propylene glycol mixtures. Fluid Phase Equilib. 2007, 262, 23−31. (12) Xu, H.-N.; Ma, S.-F. Effect of Temperature on Viscosity of Aqueous Triton X-114 Solution. J. Chem. Eng. Data 2011, 56, 2677− 2679. (13) Sun, H.; Li, M.; Jia, J.; Tang, F.; Duan, E. Measurement and Correlation of the Solubility of 2,6-Diaminohexanoic Acid Hydrochloride in Aqueous Methanol and Aqueous Ethanol Mixtures. J. Chem. Eng. Data 2012, 57, 1463−1467. (14) Sardari, F.; Jouyban, A. Solubility of Nifedipine in Ethanol + Water and Propylene Glycol + Water Mixtures at (293.2 to 313.2) K. Ind. Eng. Chem. Res. 2013, 52, 14353−14358. (15) Ben-Naim, A. Preferential solvation in two-component systems. J. Phys. Chem. 1989, 93, 3809−3813. (16) El-Sonbati, A. Z.; El-Bindary, A. A. A.; Shoair, A. G. F.; Younes, R. M. Stereochemistry of new nitrogen containing heterocyclic aldehyde. Potentiometric, Conductometric and thermodynamic studies of novel Quinoline Azodyes and their metal complexes with some transition metals. Chem. Pharm. Bull. 2001, 49, 1308−1313. (17) El-Bindary, A. A.; El-Sonbati, A. Z.; El-Mosalamy, E. H.; Ahmed, R. M. Potentiometric and Thermodynamic Studies of Azosulfonamide Drugs. X. Chem. Pap. 2003, 57, 255−258. (18) Prausnitz, J. M.; Lichtenthaler, R. N.; Azevedo, E. G. Thermodynamics of Fluid Phase Equilibria, 3rd ed.; Prentice Hall PTR: Upper Saddle River, NJ, 1999. (19) Sousa, J. M. M. V.; Almeida, J. P. B.; Ferreira, A. G. M.; Fachada, H. C.; Fonseca, I. M. A. Solubility of HFCs in lower alcohols. Fluid Phase Equilib. 2011, 303, 115−119.

the National High Technology Research and Development Program (863 Program No.2012AA021202). Notes

The authors declare no competing financial interest.

■

NOTATIONS x1 the molar fraction solubility of L-carnitine m1 the mass of L-carnitine (g) m2, m3 the mass of different solvents (g) R gas constant, 8.3145 J·mol−1·K−1 RAD the relative average deviation RMSD the root mean square deviation T temperature (K) Tm melting temperature (K) Thm harmonic temperature (K) δE influence coefficient of ethanol δA influence coefficient of acetone XE the molar fraction of ethanol in the absence of L-carnitine XA the molar fraction of acetone in the absence of L-carnitine Ji model parameters for the Jouyban−Acree model r2 correlation coefficient γ the intercept of the linear relation of lnx1 versus (1/T − 1/ Thm) ΔHd the standard change for enthalpy of the dissolving process (kJ·mol−1) ΔSd the standard change for entropy of the dissolving process (J·mol−1·K−1) ΔGd the standard change for Gibbs free energy of the dissolving process (kJ·mol−1) rH the relative contribution of the enthalpy rS the relative contribution of the entropy

■

REFERENCES

(1) Mo, J. X.; Shi, S. J.; Zhang, Q.; Gong, T.; Sun, X.; Zhang, Z. R. Synthesis, transport and mechanism of a type I prodrug: L-carnitine ester of prednisolone. Mol. Pharmaceut. 2011, 8, 1629−40. (2) Ramsay, R. R. The carnitine acyltransferases: modulators of acylCoA-dependent reactions. Biochem. Soc. Trans. 2000, 28, 182−186. (3) Belay, B.; Esteban-Cruciani, N.; Walsh, C. A.; Kaskel, F. J. The use of levo-carnitine in children with renal disease: a review and a call for future studies. Pediatr. Nephrol. 2006, 21, 308−317. (4) Nesterov, V. V.; Kolodyazhnyi, O. I. Asymmetic synthesis of a phosphorus analog of natural L-carnitine. Russ. J. Gen. Chem. 2006, 76, 1510−1511. (5) Gustavo, D. M.; Marco, G. Solubility of Acetylsalicylic Acid in Ethanol, Acetone, Propylene Glycol, and 2-Propanol. J. Chem. Eng. Data 2008, 53, 256−258. (6) Thati, J.; Nordström, F. L.; Rasmuson, Å. C. Solubility of Benzoic Acid in Pure Solvents and Binary Mixtures. J. Chem. Eng. Data 2010, 55, 5124−5127. (7) Buchowski, H.; Ksiazczak, A.; Pietrzyk, S. Solvent activity along a saturation line and solubility of hydrogen-bonding solids. J. Phys. Chem. 1980, 84, 975−979. (8) Liu, Z.; Yin, Q.; Zhang, X.; Zhang, H.; Gong, J.; Wang, J. Measurement and correlation of the solubility of 4,4′-oxydianiline in different organic solvents. Fluid Phase Equilib. 2013, 356, 38−45. (9) Krug, R. R.; Hunterlib, W. G.; Grieger, R. A. Grlegerl, Enthalpyentropy compensation. 1. Some fundamental statistical problems associated with the analysis of van’t Hoff and Arrhenius data. J. Phys. Chem. 1976, 80, 2335−2341. (10) Krug, R. R.; Hunterlib, W. G.; Grieger, R. A. Enthalpy-entropy compensation. 2. Separation of the chemical from the statistical effect. J. Phys. Chem. 1976, 80, 2341−2351. 1990

dx.doi.org/10.1021/je500078n | J. Chem. Eng. Data 2014, 59, 1984−1990