Accepted Manuscript

Experimental measurement and modelling of solubility of inosine-5’-mono-phosphate disodium in pure and mixed solvents

Fengxia Zou, Wei Zhuang, Jinglan Wu, Jingwei Zhou, Qiyan Liu, Yong Chen,Jingjing Xie, Chenjie Zhu, Ting Guo, Hanjie Ying

PII: S0021-9614(14)00139-6DOI: http://dx.doi.org/10.1016/j.jct.2014.04.023Reference: YJCHT 3924

To appear in: J. Chem. Thermodynamics

Received Date: 22 January 2014Revised Date: 23 April 2014Accepted Date: 27 April 2014

Please cite this article as: F. Zou, W. Zhuang, J. Wu, J. Zhou, Q. Liu, Y. Chen, J. Xie, C. Zhu, T. Guo, H. Ying,Experimental measurement and modelling of solubility of inosine-5’-monophosphate disodium in pure and mixedsolvents, J. Chem. Thermodynamics (2014), doi: http://dx.doi.org/10.1016/j.jct.2014.04.023

This is a PDF file of an unedited manuscript that has been accepted for publication. As a service to our customerswe are providing this early version of the manuscript. The manuscript will undergo copyediting, typesetting, andreview of the resulting proof before it is published in its final form. Please note that during the production processerrors may be discovered which could affect the content, and all legal disclaimers that apply to the journal pertain.

1

Experimental measurement and modelling of solubility of

inosine-5'-monophosphate disodium in pure and mixed

solvents

Fengxia Zou a, c, 1

, Wei Zhuang a, 1

, Jinglan Wu a, Jingwei Zhou

a, Qiyan Liu

a, Yong

Chen a, Jingjing Xie

a, Chenjie Zhu

a, Ting Guo

a, Hanjie Ying

*, a, b, c

a College of Biotechnology and Pharmaceutical Engineering, Nanjing

University of Technology, Nanjing 210009, China

b State Key Laboratory of Materials-Oriented Chemical Engineering,

Nanjing210009, China

c College of Pharmacy, Nanjing University of Technology,

Nanjing210009, China

1These authors contributed equally to this work

*Correspondence: Dr. Hanjie Ying, College of Biotechnology and

Pharmaceutical Engineering, Nanjing University of Technology, Xin

mofan Road 5, Nanjing 210009, PR China.19

E-mail address: yinghanjie@njut.edu.cn

Tel.:+86 25 86990001; fax: +86 25 58139389.

2

Abstract

The solubility of biological chemicals in solvents provide important

fundamental data and is generally considered as an essential factor in the design of

crystallization processes. The equilibrium solubility data of inosine-5'-monophosphate

disodium (5′-IMPNa2) in water, methanol, ethanol, acetone, as well as in the solvent

mixtures (methanol +water, ethanol + water, acetone + water), were measured by an

isothermal method at temperatures ranging from 293.15 K to 313.15 K. The measured

data in pure and mixed solvents were then modelled using the modified Apelblat

equation, van’t Hoff equation, λh equation, ideal model and the Wilson model. The

modified Apelblat equation showed the best modelling results, and it was therefore

used to predict the mixing Gibbs energies, enthalpies, and entropies of 5′-IMPNa2in

pure and binary solvents. The positive values of the calculated partial molar Gibbs

free energies indicated the variations in the solubility trends of 5′-IMPNa2. Water and

ethanol (in the binary mixture with water) were found to be the most effective solvent

and anti-solvent, respectively.

Keywords: Inosine-5'-monophosphate disodium, equilibrium solubility, pure

solvent, binary solvent, thermodynamic model

3

1. Introduction

With the increasingly stringent quality requirements that accompany the

development of fine chemicals, the crystallization processes and solubility in various

environments have attracted widespread attention [1, 2].Especially in the cases of

pharmaceuticals, biological macromolecular foods, and food additives [3], their

solubility in various solvents is an important property that needs to be known for the

formulation of fine chemicals and designing their production and purification

processes [4].

Inosine-5'-monophosphate disodium (5'-IMPNa2), which is known to be a kind

of nucleotide derivative, is widely used in the fields of food, pharmacological, and

health products [5]. The molecular structure of 5′-IMPNa2 is shown in figure 1. Many

studies have focused on the upstream roles of 5′-IMPNa2, but few involved in its

crystallization process, including solubility, super-solubility, meta-stable zone,

primary nucleation, and so on [6].

[Figure 1 about here]

Crystallization is an important step in industrial purification processes, and it

strongly relies on accurate equilibrium solubility that varies with temperature and

solution composition. The super-saturation of a solution during crystallization has a

direct effect on the quality of the resulting crystals, therefore, the solubility is

generally considered as essential fundamental data for controlling the product quality

during crystallization processes [7]. Furthermore, it is of great commercial value and

industrial prospect to employ a purification process that yields 5′-IMPNa2 with high

purity and a beautiful crystal habit. A detailed examination of the solubility of

5′-IMPNa2 is thus required for the sake of applying it to industrial crystallization and

production.

Solubility measurement methods as well as thermodynamic modelling have been

widely reported in literature. The most common measurement methods are the

isothermal method [8], gravimetric method [9], and a dynamic method that uses laser

monitoring as the observation technique [10]. Most studies simply involve

aqueous/methanol solvents over aqueous/ethanol solvents to fit the solubility results

to the modified Apelblat and Redlich-Kister equations [11-13].Recently, the solubility

of various compounds in organic solvents was measured and modelled using the

following models: Wilson model [7, 14], NRTL model [15], and UNIQUAC model

[16, 17], all of which relate the activity coefficients of solute with their solubility [18,

19].

In this work, the solubility of 5′-IMPNa2 in pure water, methanol, ethanol, and

acetone as well as in mixed solvents including methanol-water, ethanol-water, and

acetone-water were measured by an isothermal method with in the temperature range

from 293.15K to 313.15 K. The modified Apelblat equation, van’t Hoff equation, λh

equation, ideal model, and Wilson model were used to model the solubility of

5′-IMPNa2 in various solvents. Furthermore, the mixing Gibbs energies, enthalpies,

and entropies of 5′-IMPNa2 in pure and mixed solvents were predicted.

4

2. Experimental

2.1. Materials

The 5′-IMPNa2, with mass fraction purity greater than 0.990, was prepared by

re-crystallization. The 5’-IMPNa2 studied was obtained from Nanjing Tongkai

Biological Technology Co.(China) and re-crystallized by our laboratory. It was

re-crystallized by the method called dilution crystallization by using the anti-solvent

ethanol. It is a process that requires adding an anti-solvent to the solution to achieve a

certain degree of saturation, resulting in solute precipitation. Firstly, a suitable amount

of anti-solvent should be added into the crystallizer until the nucleation occurred.

Secondly, stop adding anti-solvent, and keep stirring for at least 4 h. Subsequently

continue adding the anti –solvent till the amount of anti-solvent reaches a certain

volume. Next wash the solute with ethanol, apply suction filtration, and dry with a

dryer of 40℃ for about 4 h. Finally we determined the purity of the solute with HPLC.

The purity of 5’-IMPNa2 is calculated by the relation between

. Methanol, ethanol, and acetone with

weight fractions of 0.995, 0.997 and 0.995, respectively were supplied by Shanghai

Chemistry Reagent Co. (China) and they are of analytical reagent grade. The source

and purity of chemicals used are given in table 1.

[Table 1 about here]

2.2. Apparatus and procedure

2.2.1. Measurement of melting properties

The melting temperature, Tm, and enthalpy of fusion, ∆Hfus, of 5′-IMPNa2 were

determined with a differential scanning calorimetry (DSC, NETZSCH STA449F3,

Germany) simultaneous thermal analyser under a nitrogen atmosphere. The masses of

samples were 6.5 mg, the temperature range from 303 K to 1073 K and the heating

rate was 30 K·min-1. The calibration of temperature and heat flow rate for DSC can be

finished at the same time during the prototype test. By testing the melting point of

different standard materials to get the ( is the difference between the

measured melting point and theoretical melting point) at different temperatures

(T).Then by fitting the curve of △T versus T, we can get a calibration curve of

temperature. The calibration of heat flow (mW) is the subsequent work of calibration

of sensitivity ( , which is obtained by fitting the curve of )

versus T. The most commonly used standard materials for DSC calibration are as

follows: In, Bi, Zn, Al and Au. The mass fraction purity of the standard materials is

at least 0.9999. The uncertainty for temperature was estimated to 0.2 K, and

uncertainty for enthalpy of fusion was estimated to be 2 %.

5

2.2.2. Solubility measurement

The solubility of 5′-IMPNa2 in pure solvents as well as mixed solvents was

measured by an isothermal method using an upright flask with magnetic stirring. An

excess amount of 5′-IMPNa2 was added into 50 mL solution at an uncertainty of

± 0.1 mg as measured by an electronic balance (BS-124S, 100 Sartorius, Germany).

The experiments were implemented in water bath under magnetic stirring (type

HH-S6, Changzhou Pacific Automation Technique Co., Ltd., China) while

maintaining different temperatures within ± 0.05 K. The mixtures were continuously

stirred for at least 24 h to obtain the solute–solvent equilibrium. The equilibrium time

was determined by measuring the concentration of 5′-IMPNa2 every 45 min until a

constant concentration was obtained. Subsequently, stirring was stopped and keep still

the upper layer was clear.

After that, a precise sampling needle was used to remove multiple samples of the

supernatant through a filter membrane to measure the volume of 1 mL with the

uncertainty of ± 0.1 mL. During the filtration process, the ambient temperature was

kept consistent with the experimental temperature. The mass of the solution was

determined by the electronic balance with an uncertainty of ± 0.1 mg. The

concentrations of 5′-IMPNa2 with an uncertainty of ± 0.01g·L-1

, were measured by

high performance liquid chromatography (HPLC; Agilent 1100, USA) using a

SepaxHP-C18 column (4.6 mm × 250 mm, 5 µm, Sepax (Jiangsu) Technologies, Inc.,

Changzhou, China). The mobile phase was 0.6 % (v/v) phosphoric acid. The column

temperature was 300.15 K, and the flow rate was 1.0 mL·min-1. Each experiment was

carried out at least three times, and then to get mean values. The estimated uncertainty

of the solubility values based on the error analysis was less than 1 %.

The solubility ( 1x ) of 5′-IMPNa2 was calculated according to Eq. (1):

∑ =

=n

i ii Mm

Mmx

1

111

/

/, (1)

where m1represents the mass of the solute (5′-IMPNa2); mi (i >1) represents the masses

of the solvents (water, methanol, ethanol and acetone); M1 represents the molar mass

of the solute, and Mi (i > 1) represents the molar mass of the solvents. The maximum

value of n is 3, where n= 1 represents the solute, n= 2 represents a pure organic

solvent, and n= 3 represents a binary solvent mixture. When n=3, 1 represents the

solute, 2 is water, 3 represents the solvent (methanol, ethanol, and acetone).

3. Modelling of solubility data

In this work, the following models were used to correlate the solubility of

5′-IMPNa2 in various solvents: the modified Apelblat equation, van’t Hoff equation,

λh equation, the ideal model and Wilson model.

6

3.1 Modified Apelblat Equation

The following modified Apelblat equation was used to correlate the solubility of

5′-IMPNa2 in various solvents [12]:

TCT

BAx lnln 1 ++= , (2)

where x1 is the solubility of 5′-IMPNa2 in the solvent in mole fraction; T is the

absolute temperature, K; A, B, and C are the model parameters. The values of A and B

correspond to the variation in solution activity coefficients, and C denotes the effect

of temperature on enthalpy of fusion [12].

3.2 van’t Hoff Equation

The van’t Hoff equation reflects the relationship between the solubility of a

solute in mole fraction and the temperature in a real solution, which is expressed as

follows [20]:

R

S

RT

Hx ddissodisso lnln

1ln∆

+∆

−= , (3)

where R is the gas constant, and ∆dissolnH and ∆dissolnS are, respectively, the dissolution

enthalpy and entropy.

3.3λh Equation

The solubility of 5′-IMPNa2 in various solvents was also modelled by the λh

equation, which is given as follows [21]:

)11

()1

1ln(1

1

mTTh

x

x−=

−+ λλ , (4)

where λ and h are two model parameters, and Tm is the melting temperature of

5′-IMPNa2. The value of λ reflects the non-ideal nature of the solution system,

whereas h denotes the enthalpy of the solution.

3.4 Ideal Model

The ideal model is a universal equation for solute–solvent equilibrium based on

thermodynamic principles. The simplified equation can be written as follows [22]:

)11

(lnln

1TTR

Hx

m

disso−

∆=γ (5)

Assuming the solution is an ideal solution (γ = 1), then Eq. (5) can be expressed

as follows:

bT

ax +=1ln , (6)

7

where a and b are the model parameters, andx1 is the solubility in mole fraction at

system temperature T.

3.5 Wilson Model

At a given temperature and pressure, the fugacity of components in the liquid

and solid phases should be the same as the phase equilibrium [14, 23].

),(),,( 1

1 PTfxPTfsl = (7)

In the solute–solvent system, the fugacity of the solute in the liquid phase can be

expressed by the following equation:

),,(),(),,( 111111 xPTfPTfxPTx sl =γ (8)

On the basis of the activity coefficient, the equilibrium solubility of the solute

may be expressed by the following simplified equation [24]:

1

1

1 ln)11

(ln γ−−∆

=TTR

Hx

m

fus (9)

where 1fus

H∆ is the enthalpy of fusion, Tm is the melting temperature, and γi is

the activity coefficient of the solute. The Wilson model, a well-established activity

coefficient model, was therefore used for the calculations in this study.

The Wilson model can be expressed in the following binary form [25]:

)()ln(ln1212

21

2121

12221211

xxxxxxx

Λ+

Λ−

Λ+

Λ+Λ+−=γ , (10)

where

)exp( 2112

1

212

RT

λλ

ν

ν −−=Λ (11)

and

)exp( 1221

2

121

RT

λλ

ν

ν −−=Λ (12)

Here, ∆λ12and ∆λ21 are the cross-interaction energy parameters that can be fitted using

experimental solubility, and ν1 and ν2 are the molar volumes of the solute and solvent,

respectively. For pure organic solvents, x2 can be simplified to be 1. The final

equation can then be expressed as:

112

21

121

121211

1)ln(ln

xxx

Λ+

Λ−

Λ+

Λ+Λ+−=γ (13)

The parameters of the five models are given in the supporting information. The

8

average absolute deviation (AAD) and average relative deviation (ARD) were used to

identify the differences between the calculated and measured results, and they are

calculated according to Eq. (14) and (15):

∑ =

−=

N

ii

cal

ii

x

xx

NAAD

1,1

,1,11 (14)

1 1

1

calx x

ARDx

−= , (15)

where N is the number of experimental points, and x1,i and x1,ical

are the experimental

and calculated solubility, respectively.

3.6 Thermodynamic properties of pure components

Figure 2 shows the thermal analysis results of 5′-IMPNa2. The DSC results

indicate that the melting temperature, Tm, and enthalpy of fusion, ∆fusH of 5′-IMPNa2

are 199.08 oC and 17.39 kJ·mol

-1, respectively. The entropy of fusion, ∆fusS of

5′-IMPNa2 was calculated using the following equation:

m

fus

fusT

HS

∆=∆ (16)

[Figure 2 about here]

The thermodynamic properties of the pure component are useful for modelling the

solubility of 5′-IMPNa2with the Wilson model.

4. Results and discussion

4.1. Measured solubility of 5′-IMPNa2 and modelling results

[Table 2 about here]

The molar volumes of 5′-IMPNa2 and the solvents (water, methanol, ethanol, and

acetone), shown in table 1, were used for modelling with the Wilson model. The mole

volume of 5’-IMPNa2 was calculated from the density of 5’-IMPNa2, 1.536 g/cm3,

which was measured using a pycnometer method [26] and the molar volumes of

solvents were the literature values [27]. Figure 3 shows the solubility of 5′-IMPNa2 in

four different pure solvents. As shown in figure 3, the solubility of 5′-IMPNa2 in water,

methanol, ethanol, and acetone are dependent on temperature, and the solubility of

5′-IMPNa2are increased with an increase in temperature. At any given temperature,

the solubility of 5′-IMPNa2 in the various pure solvents follows the order of

water>methanol>ethanol>acetone. From the molecular structure of 5′-IMPNa2 as

shown in Figure 1, it has a phosphate group，a purine ring, and a large number of

9

hydroxyl moieties. According to the Mullin’s theory [28], based on the nature of their

intermolecular bonding interactions, solvents may be conveniently divided into three

main categories: polar protic (such as water, methanol, and ethanol), dipolar

aprotic(such as furfural and nitrobenzene), and non-polar aprotic (such as hexane,

benzene, and acetone). 5′-IMPNa2 exhibits polar molecular bonding. Therefore, it can

be easily dissolved in polar protic solvents because its molecular bonding has

sufficient polarity to break and replace hydrogen bonds. So, for the polar protic

solvents; in this case water, methanol, and ethanol, the solubility of 5′-IMPNa2 in

them is consistent with their polarities. However, in the case of the non-polar aprotic

solvent acetone, characterized by low dielectric constants, molecular interaction takes

place through the weak van der Waals forces. Dipolar and polar protic solutes are

generally found to have very low solubility in such non-polar aprotic solvents because

the van der Waals forces cannot break and replace the hydrogen bonds to ensure

solute dissolution. To sum up, the solubility behaviour of 5′-IMPNa2 in various

solvents is in accordance with the rule “like dissolves like” [28].According to the

theory of Bronsted-Lowry acid and alkali, which can provide protons are exactly acid

substances, which means, the acid strength is consistent with solubility data. Based on

the organic chemistry knowledge, the four solvents of acid strength rank as: water>

methanol>ethanol>acetone. This theory is applicable for the consequence in this

article, that is to say, if the solute belongs to polar protic, just to compare the ability of

the solvents to provide the hydrogen atom. Methanol, ethanol belong to polar protic,

and acetone is non-polar protic, the ability to provide the hydrogen atom arrange as:

methanol> ethanol> acetone. The result acquired from the explanation of theory is the

same with the rule “like dissolves like”.

[Figure 3 about here]

The parameters and AAD% of the models are presented in tables 3 and 4.

According to table 3, compared with the other four models, the AAD% from the

modified Apelblat equation is the lowest, which indicates that the modified Apelblat

equation gives the best fitting results of the solubility data. The AAD% from fitting

data using the modified Apelblat equation of the four solvents are as follows: 0.40

(water), 0.84(methanol), 1.34(ethanol), and 1.70(acetone).Therefore, all data were

fitted using this equation in our study. The fitted parameters from table 4 are quite

accurate.

[Table 3 about here]

[Table4 about here]

4.2 Prediction of dissolution enthalpy, entropy, and the molar Gibbs energy in pure

organic solvent

According to the van’t Hoff equation [29], the apparent standard molar enthalpy

10

change of solution could be related to the temperature and the solubility of the solute,

as represented by following equation:

)1

(

lnln

T

x

R

Ho

so

∂

∂−=

∆. (17)

Over a limited temperature interval (293.15–313.15) K, the change in heat

capacity of a solution may be assumed to be constant. Hence, the values of ∆solnHº

would be valid for the mean temperature (303.15 K) [30].

Therefore, on combining Eq. (17) with Eq. (2), the ∆solnH can be calculated using

the resulting Eq. (18) and Eq. (19).

)()1

ln 1ln ( mean

o

so cTbR

T

xRH −−=

∂

∂−=∆ . (18)

and

o

somean

o

somean

o

so STHRTG lnlnln intercept ∆−∆=×−=∆ (19)

in which, the intercept can be obtained by plotting ln 1x versus (1/T -1/Tmean.) (as

shown in figure 4).

[Figure 4 about here]

The Gibbs energy of solution can be calculated using the following equation:

o

somean

o

so

o

so STHG lnlnln ∆−∆=∆ (20)

The solution process may be represented by the following hypothetic stages:

ution)Solute(sol→uid)Solute(liq→id)Solute(sol (21)

Where the respective partial processes toward the solution are solute fusion and

mixing at the same temperature, which allows calculate the partial thermodynamic

contributions to the overall solution process by means of the following equations,

respectively [31].

∆solnHo = ∆fusH + ∆mixH

o (22)

∆solnSo = ∆fusS + ∆mixS

o (23)

∆solnGo = ∆fusG + ∆mixG

o (24)

The solute is dried before the experiment, and is cooled to room temperature in a

drier. The temperature and humidity of the experimental environment is invariable.

Based this, the possibility of the solute melting is very low, so we ignored the partial

11

contribution for practical purposes in this analysis. So the ∆solnG can be thought to be

equal to ∆mixGo approximately, the same with enthalpy and entropy.

The calculated dissolution enthalpy, entropy, and Gibbs energy are presented in

table 5. Results show that the o

so Hln∆ of 5′-IMPNa2 in each solvent in the

experimental temperature range is positive ( o

so Hln∆ > 0), indicating dissolution is

endothermic, which explains the increasing solubility of 5′-IMPNa2 with increasing

temperature.

[Table 5 about here]

Dissolution is an endothermic process because the interactions between the

5′-IMPNa2 molecules and the solvent molecules are more powerful than those

between the solvent molecules themselves. The positive o

so Hln∆ and o

so Sln∆ values

in water, methanol, ethanol, and acetone indicate that the dissolution process of

5′-IMPNa2 in these four pure solvents is all entropy-driven [32].

As can be seen from the data presented in the tables, the values of the Gibbs

energy of solution are all positive, and decrease with the increasing temperatures for

all the solvents. Lower ∆solnGo values correspond to higher solubility and more

favourable dissolution. In addition, the appropriate positive ∆solnGo values for all

solutions indicate that 5′-IMPNa2 can be easily crystallized from any of the studied

solvents.

4.3. Solubility of 5′-IMPNa2 in binary solvents

The solubility data laid the foundation for the later study on the solute

crystallization process. It is a process that requires adding an anti-solvent to the

solution to achieve a certain degree of saturation, resulting in solute precipitation.

Based on the solubility data in the pure organic solvents—water, methanol, ethanol,

and acetone—it can be seen that 5′-IMPNa2 is soluble in all of them to varying

degrees. However, these organic solvents and water are not appropriate for industrial

production in their pure states. Therefore, another solvent needs to be introduced into

an aqueous solution system to reduce the saturation in the mother liquid, which is

water. The significance of measuring solubility in binary mixtures is to choose the

best anti-solvent for industrial production considering the cost and safety of the

anti-solvent. So (methanol + water), (ethanol + water), and (acetone+ water) in

different ratios were investigated across different temperature ranges in our study.

During the crystallization process, we need an initial amount ethanol to acquire

12

saturation for nucleation and growth. A large number of experiments proof that the

best ethanol mole fraction for nucleation is from 0.15 to 0.3 for different initial

concentrations, so we investigated the solubility of the water mole fraction from 0.0 to

0.5. The solubility data of pure organic solvent were used to select a best anti-solvent,

so we did not compare them to the solubility with different mole fraction in the binary

part. The basic results obtained are listed in table 6 in detail, and the parameters fitted

by the modified Apelblat equation are shown in table 7. Based on the AAD% of the

different binary mixtures, the model seems to fit the results fairly well.

[Table 6 about here]

It can be observed from figures 5a, 5b, 5c that the solubility of 5′-IMPNa2 in the

different binary solvents at a given temperature ranks as follows: (methanol +

water) > (ethanol + water) > (acetone + water), which are in great accordance with the

solubility in pure solvents.

[Table 7 about here]

In addition, the solubility values increased with the increasing mole fraction of

water at any given temperature. This is due to the physical properties of 5′-IMPNa2

and water [31]. The polar molecules of water are able to solvate the Na+

and IMP2-

ions because their partially-charged portions can orientate appropriately towards the

ions in response to electrostatic attraction. This solvation stabilizes the system and

creates a hydration shell [33, 34]. In addition, the amino group and hydroxyl groups in

the IMP2-

ions to associate with water via hydrogen bonding. Hence, 5′- IMPNa2

possesses high solubility in water.

[Figure 5 about here]

However, in binary solvents, the hydroxyl group from water associates with the

alcohols and acetone via hydrogen bonding more easily than it associates

with5′-IMPNa2. As a result, the addition of organic solvents decreases the number of

water molecules available for solvation of 5′- IMPNa2, leading to the lower solubility

of 5′-IMPNa2. Consequently, the organic solvents demonstrated an anti-solvent effect

for 5′- IMPNa2. The solubility of 5′- IMPNa2 exhibited the largest change in the

(ethanol + water) system, indicating that ethanol is an effective anti-solvent for

5′-IMPNa2 that meets the requirements of food and drug processing. He reasons why

ethanol is an effective anti-solvent are as follows: 1) According to the value of

in pure solvents and in binary solvents is the lowest in ethanol, which means during

the dissolution process the less energy is required to overcome the cohesive forces

13

with the solute and the solvent. 2) As is known to us all, 5’-IMPNa2 is a kind of food

additive. The residue of solvent in final products has strict standards. So among the

three anti-solvents, toxicity of ethanol is the smallest. 3) In view of the industrial cost,

ethanol is much cheaper than methanol and acetone.

4.4. The thermodynamic properties of the binary mixture solutions

Based on Eqs.17 to 19, the thermodynamic properties were calculated for binary

solutions as well as for pure organic solvents, as shown in figure 6.The calculated

values of standard molar Gibbs energy, entropy, and enthalpy of 5′-IMPNa2 are listed

in table 8.

[Figure 6 about here]

[Table 8 about here]

The positive values of o

so Hln∆ and o

so Sln∆ reveal that the dissolution of

5′-IMPNa2 in the mixtures is an entropy-driven process. It can be also seen that the

values of o

so Hln∆ decreased with increasing mole fraction of water, and attained a

minimum at 2x =0.6082 in the (methanol + water) mixture, a trend reflected by

entropy o

so Sln∆ too. Similarly, both o

so Hln∆ and o

so Sln∆ exhibited minimum values

at the same point for all binary mixtures ( 2x = 0.5990 in ethanol + water and 2x =

0.6038 in acetone + water). In pure solvents, the value of o

so Hln∆ in methanol is

higher than those in acetone, water, and ethanol. And in binary solvents too, the value

of o

so Hln∆ for the mixture containing ethanol is the lowest. The total value of the

enthalpy for the solution process includes the several kinds of interactions that occur

during dissolution. Therefore, a higher value of total enthalpy indicates that more

energy is required to overcome the cohesive forces within the solute and the solvent

during the dissolution process, which also signifies the stronger dependence of

solubility on temperature [35]. The results indicate that ethanol is the best solvent for

14

the crystallization process.

The positive values of the changes in Gibbs energy of solution indicate that the

process was not spontaneous. Eq.(25) and (26) were used to compare the relative

contributions of the enthalpy and entropy of the solution process to the Gibbs

energy[36].

o

soanm

o

so

o

so

HSTH

H

lneln

ln100%

∆+∆

∆=ζ . (25)

o

somean

o

so

o

somean

TSSTH

ST

lnln

ln100%

∆+∆

∆=ζ . (26)

The values of Hζ% and TSζ% in binary mixtures were calculated and are listed

in table 7. The value of Hζ% exceeds 60, which indicates that the main contributing

factor to the standard Gibbs energy was the enthalpy of dissolution of 5′-IMPNa2 in

all the mixtures studied [36].

5. Conclusions

The solubility of 5′-IMPNa2 in various pure solvents and binary mixtures were studied

over the temperature range of (293.15–313.15) K. From the results, the solubility in

the pure solvents studied at any given temperature could be ranked as follows:

water > methanol > ethanol > acetone. Also it can be seen that the solubility of

5′-IMPNa2 increases with increasing temperature at a constant solvent composition

within the temperature range studied, and increases with the increasing water ratio

over the same temperature range for binary solvent mixtures. All the thermodynamic

properties associated with the dissolution of 5′-IMPNa2 in binary mixtures and pure

solvent are positive, which indicate that dissolution of 5′-IMPNa2 in the selected

solvents is not spontaneous and is an endothermic process. The experimental

solubility values in solvent mixtures were correlated based on the modified Apelblat

equation because it gave the best fitting results for the pure organic solvents. The

modified Apelblat equation was also used to study the mixing properties of binary

15

mixtures, and the final results are found to be well correlated with the experimental

solubility data. Water was found a better solvent for 5′-IMPNa2 than the others. In

contrast, ethanol and methanol could be used as effective anti-solvents in the

crystallization process, with the former being more effective as an anti-solvent. The

experimental solubility results and equations presented in this study can be used to

optimize the crystallization conditions of 5′-IMPNa2 in practical applications. A

comprehensive study on the impact of pH on 5′-IMPNa2 will be conducted later.

AUTHOR INFORMATION

Corresponding Author

*Phone:+86 25 86990001. Fax: +86 25 58139389. E-mail: yinghanjie@njtech.edu.cn

Acknowledgement

This work was supported by the National Basic Research Program of China (No.

2011CBA00806), Natural Science Foundation of Jiangsu Grants (No. BK20130929,

BK2011031), Jiangsu Postdoctoral Science Foundation (1301038B), National

Outstanding Youth Foundation of China (No.:21025625), National High-Tech

Research and Development Plan of China (2012AA021202), Natural Science

Foundation of China Grants (No. 21106070).

Reference

[1] B.Y. Shekunov, P. York, Crystallization processes in pharmaceutical technology and drug delivery

design, Journal of crystal growth, 211 (2000) 122-136.

[2] W.J. Habraken, J. Tao, L.J. Brylka, H. Friedrich, L. Bertinetti, A.S. Schenk, A. Verch, V.

Dmitrovic, P.H. Bomans, P.M. Frederik, Ion-association complexes unite classical and non-classical

theories for the biomimetic nucleation of calcium phosphate, Nature communications, 4 (2013) 1507.

[3] E. Sheikholeslamzadeh, S. Rohani, Solubility Prediction of Pharmaceutical and Chemical

Compounds in Pure and Mixed Solvents Using Predictive Models, Industrial & Engineering Chemistry

Research, 51 (2012) 464-473.

[4] R.J. Davey, S.L. Schroeder, J.H. ter Horst, Nucleation of Organic Crystals—A Molecular

Perspective, Angewandte Chemie International Edition, 52 (2013) 2166-2179.

[5] B. Rifkin, L.M. Bartoshuk, Taste synergism between monosodium glutamate and disodium

5′-guanylate, Physiology & Behavior, 24 (1980) 1169-1172.

[6] T. Kumazawa, M. Nakamura, K. Kurihara, Canine taste nerve responses to umami substances,

Physiology & behavior, 49 (1991) 875-881.

[7] J. Li, Z. Wang, Y. Bao, J. Wang, Solid–Liquid Phase Equilibrium and Mixing Properties of

Cloxacillin Benzathine in Pure and Mixed Solvents, Industrial & Engineering Chemistry Research, 52

(2013) 3019-3026.

[8] X. Zhou, J. Fan, N. Li, Z. Du, H. Ying, J. Wu, J. Xiong, J. Bai, Solubility of l-phenylalanine in

16

water and different binary mixtures from 288.15 to 318.15K, Fluid Phase Equilibria, 316 (2012) 26-33.

[9] D. Jia, L. Wang, C. Li, X. Wang, Solid–liquid phase equilibrium of glyphosate in selected solvents,

Fluid Phase Equilibria, 327 (2012) 1-8.

[10] B.S. Liu, H. Sun, J.K. Wang, Q.X. Yin, Solubility of disodium 5′-guanylate heptahydrate in

aqueous methanol mixtures, Food Chemistry, 128 (2011) 218-221.

[11] J. Yu, T. Ma, A. Li, X. Chen, Y. Chen, J. Xie, J. Wu, H. Ying, Solubility of disodium cytidine

5′-monophosphate in different binary mixtures from 288.15 K to 313.15 K, Thermochimica Acta,

565(2013)1-7.

[12] Z. Du, Y. Hu, P. Wang, J. Zhou, J. Xiong, H. Ying, J. Bai, Conformers and hydrogen bonds in

cytidine 5′-diphosphocholine sodium single crystals grown from a mixture of ethanol and water,

Journal of Molecular Structure, 985 (2011) 227-232.

[13] B. Wang, L. Wang, J. Zhang, W. Li, J. Wang, Measurement and correlation for solubility of

dexibuprofen in different solvents from 263.15 to 293.15K, Thermochimica Acta, 540 (2012) 91-97.

[14] J. Chen, Z.X. Zeng, W.L. Xue, D. Wang, Y. Huang, Determination and Correlation of Solubility

of Decahydropyrazino pyrazine in Methanol, Ethanol, and 2-Propanol, Industrial & Engineering

Chemistry Research, 50 (2011) 11755-11762.

[15] W. Omar, J. Ulrich, Solid liquid equilibrium, metastable zone, and nucleation parameters of the

oxalic acid-water system, Crystal growth & design, 6 (2006) 1927-1930.

[16] J. Lu, Q. Lin, Z. Li, S. Rohani, Solubility of l-phenylalanine anhydrous and monohydrate forms:

Experimental measurements and predictions, Journal of Chemical & Engineering Data, 57 (2012)

1492-1498.

[17] P. Oracz, M. Góral, Application of the Unified Functional Activity Coefficient (UNIFAC) and

Analytical Solution of Groups (ASOG) for the calculation of mutual solubilities in water systems of

alkanes, arenes, and alkanols, Journal of Chemical & Engineering Data, 56 (2011) 4853-4861.

[18] M. Tao, Z. Wang, J. Gong, H. Hao, J. Wang, Determination of the Solubility, Dissolution

Enthalpy, and Entropy of Pioglitazone Hydrochloride (Form II) in Different Pure Solvents, Industrial &

Engineering Chemistry Research, 52 (2013) 3036-3041.

[19] H. Yan, R. Li, Q. Li, J. Wang, J. Gong, Solubility of Minoxidil in Methanol, Ethanol, 1-Propanol,

2-Propanol, 1-Butanol, and Water from (278.15 to 333.15) K, Journal of Chemical & Engineering Data,

56 (2011) 2720-2722.

[20] N. Aldabaibeh, M.J. Jones, A.S. Myerson, J. Ulrich, The solubility of orthorhombic lysozyme

crystals obtained at high pH, Crystal Growth and Design, 9 (2009) 3313-3317.

[21] D.Q. Li, D.Z. Liu, F.-A. Wang, Solubility of 4-methylbenzoic acid between 288 K and 370 K,

Journal of Chemical & Engineering Data, 46 (2001) 234-236.

[22] Z. MAO, X. SUN, X. LUAN, Y. WANG, G. LIU, Measurement and correlation of solubilities of

adipic acid in different solvents, Chinese Journal of Chemical Engineering, 17 (2009) 473-477.

[23] M.S. Selvan, M.D. McKinley, R.H. Dubois, J.L. Atwood, Liquid-liquid equilibria for toluene+

heptane+ 1-ethyl-3-methylimidazolium triiodide and toluene+ heptane+ 1-butyl-3-methylimidazolium

triiodide, Journal of Chemical & Engineering Data, 45 (2000) 841-845.

[24] H. Renon, J.M. Prausnitz, Local compositions in thermodynamic excess functions for liquid

mixtures, AIChE journal, 14 (1968) 135-144.

[25] M. Mirmehrabi, S. Rohani, L. Perry, Thermodynamic modelling of activity coefficient and

prediction of solubility: Part 2. Semipredictive or semiempirical models, Journal of pharmaceutical

sciences, 95 (2006) 798-809.

17

[26] H. Blanco-Canqui, R. Lal, W.M. Post, R. Izaurralde, M. Shipitalo, Organic carbon influences on

soil particle density and rheological properties, Soil Science Society of America Journal, 70 (2006)

1407-1414.

[27] J. Gmehling, U. Onken, W. Arlt, Vapor–Liquid Equilibrium Data Collection; DECHEMA:

Frankfurt, Germany, 1977.

[28] J.W. Mullin, Crystallization, Butterworth-Heinemann, Oxford, 2001.

[29] J.B. Chaires, Possible origin of differences between van't Hoff and calorimetric enthalpy estimates,

Biophysical chemistry, 64 (1997) 15-23.

[30] M.A. Ruidiaz, D.R. Delgado, F. Martínez, Y. Marcus, Solubility and preferential solvation of

indomethacin in 1, 4-dioxane+ water solvent mixtures, Fluid Phase Equilibria, 299 (2010) 259-265.

[31] M. Gantiva, F. Martínez, Thermodynamic analysis of the solubility of ketoprofen in some

propylene glycol+water cosolvent mixtures, Fluid Phase Equilibria, 293 (2010) 242-250.

[32] N. Wang, Q. Fu, G. Yang, Determination of the solubility, dissolution enthalpy and entropy of

icariin in water, ethanol, and methanol, Fluid Phase Equilibria, 324 (2012) 41-43.

[33]C. Faralli, M. Pagliai, G. Cardini, V. Schettino, The solvation dynamics of Na+ and K+ ions in

liquid methanol, Theoretical Chemistry Accounts, 118 (2007) 417-423.

[34] R.J. Wheatley, The solvation of sodium ions in water clusters: intermolecular potentials for

Na+-H2O and H2O-H2O, Molecular Physics, 87 (1996) 1083-1116.

[35] B. Schröder, L.M. Santos, I.M. Marrucho, J.A. Coutinho, Prediction of aqueous solubilities of

solid carboxylic acids with COSMO-RS, Fluid Phase Equilibria, 289 (2010) 140-147.

[36] G.L. Perlovich, S.V. Kurkov, A.N. Kinchin, A. Bauer-Brandl, Thermodynamics of solutions III:

Comparison of the solvation of (+)-naproxen with other NSAIDs, European journal of pharmaceutics

and biopharmaceutics, 57 (2004) 411-420.

18

Table 1

Source and purity of the chemicals used in this work

Table2

Molar volume ( )mV of pure components

Table3

The solubility(x1) of 5′-IMPNa2 in different pure solvents at different temperatures (T),

pressures p = 0.1 MPa and AAD%s from the different models

Table 4

Parameters of models to predict the solubility data of 5′-IMPNa2 in different pure

solvents

Table 5

Predictions of the thermodynamic parameters in the various pure solvents and the

Gibbs energy of solution.

Table6

The solubility(x1) of 5′-IMPNa2with different x2in binary solvent mixtures at different

temperatures and pressure p = 0.1 MPa

Table 7

Parameters of the modified Apelblat equation used to obtain solubility data for the

binary mixtures

Table 8

Thermodynamic functions related to the dissolution process of 5′-IMPNa2 in binary

19

mixtures containing organic solvents ( 2x ) at T= 303.15 K.

Figure 1

Molecular structure of 5′-IMPNa2

Figure 2

DSC curve of 5′-IMPNa2 with a melting temperature of 199.08 ℃ and enthalpy of

fusion of 17.39 kJ·mol-1

Figure 3

Experimentally measured solubility (x1)of 5′-IMPNa2(x1), in different pure solvents：

（■）water；（●）methanol；（▲）ethanol; (★)acetone. The solid lines describe the

calculated results using the modified Apelblat equation

Figure 4

Temperature dependence of the solubility(x1) of 5′-IMPNa2 in pure solvents:（■）

water；（●●）methanol；（▲）ethanol; (★) acetone. The solid lines fitted by the liner

line to get the intercept for Eq.19.

Figure 5

Solubility(x1) of 5′-IMPNa2 in the water + organic solvent mixture at different

temperatures with different water mole fraction(x2): (◆)T = 313.15K; (▼) 308.15K;

(▲) 303.15K; (●) 298.15K; (★) 293.15K. ( a )water + methanol; ( b )water +ethanol;

( c ) water + acetone.

Figure 6

Temperature dependence of solubility(x1) of 5′-IMPNa2 in different binary system

with different water mole fraction(x2):

a. water + methanol (★)x2=0.9049; (●)x2=0.8017; (▲)x2=0.7044; (▼)x2=0.6082;

(◆)x2=0.5099 ;

20

b. water + ethanol

(★)x2=0.8985;(●)x2=0.7984;(▲)x2=0.6984;(▼)x2=0.5990;(◆)x2=0.4999;

c water + acetone(★)x2=0.9006;(●)x2=8028;(▲)x2=0.7030;(▼)x2=6038;(◆)x2=5046.

21

Figure 1 Molecular structure of 5′-IMPNa2

22

Figure 2 DSC curve of 5′-IMPNa2 with a melting temperature of 199.08 ℃ and

enthalpy of fusion of 17.39 kJ·mol-1

23

Figure 3 Experimentally measured solubility (x1)of 5′-IMPNa2(x1), in different pure

solvents：（■）water；（●）methanol；（▲）ethanol; (★)acetone. The solid lines

describe the calculated results using the modified Apelblat equation

24

Figure 4 Temperature dependence of the solubility(x1) of 5′-IMPNa2 in pure

solvents: （■）water；（●●）methanol；（▲）ethanol; (★) acetone. The solid lines

fitted by the liner line to get the intercept for Eq.19.

25

26

Figure 5 Solubility(x1) of 5′-IMPNa2 in the water + organic solvent mixture at

different temperatures with different water mole fraction(x2): (◆) T = 313.15K; (▼)

308.15K; (▲) 303.15K; (●) 298.15K; (★) 293.15K. ( a )water + methanol; ( b )water

+ethanol; ( c ) water + acetone.

27

28

Figure 6 Temperature dependence of solubility(x1) of 5′-IMPNa2 in different binary system with

different water mole fraction(x2).

a. water + methanol (★)x2=0.9049;(●)x2=0.8017;(▲)x2=0.7044;(▼)x2=0.6082;(◆)x2=0.5099 ;

b. water + ethanol (★)x2=0.8985;(●)x2=0.7984;(▲)x2=0.6984;(▼)x2=0.5990;(◆)x2=0.4999;

c water + acetone(★)x2=0.9006;(●)x2=8028;(▲)x2=0.7030;(▼)x2=6038;(◆)x2=5046.

29

Table 1 Source and purity of the chemicals used in this work.

Compounds Mass fraction purity by

HPLC Source

Method of

purification

5’-IMPNa2

Nanjing Tongkai

Biological Technology

Co. (China)

Re-crystallization

Methanol Shanghai Chemistry

Reagent Co. (China)

Ethanol Shanghai Chemistry

Reagent Co. (China)

No further

purification

Acetone Shanghai Chemistry

Reagent Co. (China)

Water De-ionised water Prepared in laboratory

30

Table 2 Molar volume ( )mV of pure components

5′-IMPNa2 Water Methanol Ethanol Acetone

mV /(cm3·mol

-1) 255.32 18.00 40.45 58.32 73.71

(1) mV is the molar Volume. ρρ

M

Mm

m

nVVm === .

( 2 ) The mV of methanol, ethanol, and acetone are from the reference [27].

31

Table 3 The solubility(x1) of 5′-IMPNa2 in different pure solvents at different

temperatures (T), pressure p = 0.1 MPa and AAD% from the different models

T/K 103x1

111 /)( xxxcal−

Modified

Apelblat λ h

Van’t

Hoff

Ideal

model

Wilson

model

Water

293.15 13.34 -0.3282 -0.0278 0.1658 0.1658 0.0407

298.15 15.52 0.8551 0.5792 0.5908 0.5908 0.0264

303.15 17.59 -0.5958 -0.9962 -1.096 -1.096 -0.0101

308.15 20.34 -0.0930 -0.1883 -0.3220 -0.3220 -0.0301

313.15 23.47 0.1482 0.7388 0.6497 0.6497 -0.0531

AAD% 0.4041 0.5061 0.5648 0.5648 3.210

Methanol

293.15 2.196 -0.4814 -1.589 -3.759 -3.759 -0.4553

298.15 3.118 0.7929 3.116 2.508 2.508 -0.2193

303.15 4.174 0.5729 2.970 3.741 3.741 -0.0817

308.15 5.256 -1.627 -2.300 -0.1162 -0.1162 -0.0142

313.15 6.659 0.7223 -6.185 -2.581 -2.581 0.0510

AAD% 0.8393 3.232 2.541 2.541 6.431

Ethanol

293.15 0.9510 -0.2405 0.0039 0.1700 0.1700 0.0186

298.15 1.112 1.4093 1.2600 1.192 1.192 0.0175

303.15 1.224 -2.9042 -3.126 -3.327 -3.327 -0.0363

308.15 1.482 2.3837 2.402 2.199 2.199 0.0042

313.15 1.650 -0.7397 -0.216 -0.3189 -0.3189 -0.0318

AAD% 1.334 1.401 1.442 1.442 2.167

Acetone

293.15 0.1275 1.344 1.255 1.187 1.187 0.0701

298.15 0.1452 -2.877 -2.555 -2.788 -2.788 0.0021

303.15 0.1764 0.2167 0.677 0.3759 0.3759 -0.0020

308.15 0.2117 2.636 2.987 2.707 2.707 -0.0134

313.15 0.2364 -1.411 -1.395 -1.576 -1.576 -0.0898

AAD% 1.670 1.774 1.727 1.727 3.548

(1) Standard uncertainties u are u(T) = ±0.05 K , (x1) = 0.01, and (p)= 0.02

(2) Standard uncertainty u was calculated using standard deviation (SD); x1 is the mole fraction of

32

the solubility of 5’-IMPNa2 at the experimental temperature T.

(3) AAD is the average absolute deviation.

33

Table 4 Parameters of models selected to predict the solubility of

5′-IMPNa2 in different pure solvents

Model Parameter Water Methanol Ethanol Acetone

Modified

Apelblat

A -118.22 810.71 -100.05 40.671

B 2963.74 -41115.2 2041.78 -4745.79

C 18.27 -119.1 15.16 -5.891

λh λ 0.3129 2.139 0.0200 0.0049

h 7865.09 2480.75 119491 579062

van’t Hoff

o

so Hln∆ /(kJ·mol-1) 21.37 41.89 21.21 24.62

o

so Sln∆ /(J·K-1·

mol-1)

4.450 92.33 14.48 9.334

Ideal model A -2570.65 -5039.03 -2550.66 -2961.48

B 4.450 11.11 1.741 1.123

Wilson

model

12λ∆ -2381.12 3001.61 4500.31 10358.8

21λ∆ 6727.47 2307.86 25024.83 8705.74

(1) A, B and C are parameters of Apelblat equation

(2) λ and h are parameters of hλ equation.

(3) H∆ and S∆ are parameters of Van’t Hoff equation

(4) a and b are parameters of the ideal model

(5) 12λ∆ and 21λ∆ are parameters of the Wilson model.

34

Table 5 Predictions of the thermodynamic properties of 5′-IMPNa2 in the various

pure solvents and the Gibbs energy of the solution.

T/K 293.15 298.15 303.15 308.15 313.15

Water

o

so Hln∆ /(kJ·mol-1

) 21.4 21.4 21.4 21.4 21.4

o

so Sln∆ /(J·K-1

·mol-1

) 37.0 37.0 37.0 37.0 37.0

o

so Gln∆ /(kJ·mol-1

) 10.5 10.3 10.2 9.97 9.79

Methanol

o

so Hln∆ /(kJ·mol-1

) 41.9 41.9 41.9 41.9 41.9

o

so Sln∆ /(J·K-1

·mol-1

) 92.3 92.3 92.3 92.3 92.3

o

so Gln∆ /(kJ·mol-1

) 14.8 14.4 13.9 13.4 13.0

Ethanol

o

so Hln∆ /(kJ·mol-1

) 21.2 21.2 21.2 21.2 21.2

o

so Sln∆ /(J·K-1

·mol-1

) 14.5 14.5 14.5 14.5 14.5

o

so Gln∆ /(kJ·mol-1

) 17.0 16.9 16.8 16.7 16.7

Acetone

o

so Hln∆ /(kJ·mol-1

) 24.6 24.6 24.6 24.6 24.6

o

so Sln∆ /(J·K-1

·mol-1

) 9.34 9.34 9.34 9.34 9.34

o

so Gln∆ /(kJ·mol-1

) 21.9 21.8 21.8 21.7 21.7

（1） o

so Hln∆ = solution (mixing) enthalpy, o

so Sln∆ = solution (mixing) entropy, o

so Gln∆ =

solution (mixing) Gibbs energy.

35

Table 6 The solubility(x1) of 5′-IMPNa2 with different x2 in binary solvent mixtures

at different values of temperature and pressure p = 0.1 MPa

2x 1

310 x ARD%

Methanol + water (T=293.15K)

0.5099 7.650 0.3491

0.6082 8.179 -0.7589

0.7044 8.631 0.1139

0.8017 9.332 0.6292

0.9049 9.989 -0.3447

Methanol + water (T=298.15K)

0.5099 9.869 0.2077

0.6082 10.48 -0.7666

0.7044 11.42 1.030

0.8017 11.97 -0.6163

0.9049 12.73 0.1331

Methanol + water (T=303.15K)

0.5099 12.39 -0.2488

0.6082 13.19 0.3526

0.7044 14.39 0.4836

0.8017 15.16 -1.032

0.9049 16.02 0.4407

Methanol + water (T=308.15K)

0.5099 15.40 0.2824

0.6082 16.37 -0.4778

0.7044 17.10 -0.3322

0.8017 17.88 0.9392

0.9049 19.51 -0.4285

Methanol + water (T=313.15K)

0.5099 18.35 0.1021

0.6082 19.30 -0.0662

0.7044 20.13 -0.4141

0.8017 20.84 0.6385

0.9049 22.77 -0.2516

Ethanol + water (T=293.15K)

0.4999 6.719 0.2998

0.5990 7.276 -1.063

0.6984 7.703 1.375

0.7984 8.805 -0.7765

0.8985 9.731 0.1582

Ethanol + water (T=298.15K)

36

0.4999 7.929 -0.0480

0.5990 8.980 0.0417

0.6984 9.617 0.1826

0.7984 10.78 -0.2898

0.8985 12.18 0.1176

Ethanol + water (T=303.15K)

0.4999 10.01 -0.0355

0.5990 10.87 0.2035

0.6984 11.69 -0.4050

0.7984 13.01 0.3368

0.8985 15.33 -0.1024

Ethanol + water (T=308.15K)

0.4999 11.96 -0.1819

0.5990 13.07 0.3863

0.6984 14.14 -0.0745

0.7984 15.40 -0.3124

0.8985 17.88 0.1677

Ethanol + water (T=313.15K)

0.4999 14.15 0.9797

0.5990 15.71 -2.523

0.6984 16.64 1.439

0.7984 18.15 0.6210

0.8985 20.90 -0.5751

Acetone + water(T=293.15K)

0.5046 6.073 0.2257

0.6038 6.792 -0.4997

0.7030 7.134 0.0989

0.8028 7.506 0.3813

0.9006 8.233 -0.2146

Acetone + water(T=298.15K)

0.5046 7.926 -0.1273

0.6038 8.555 0.0835

0.7030 8.844 0.5416

0.8028 9.580 -0.8304

0.9006 10.64 0.3276

Acetone + water(T=303.15K)

0.5046 9.788 0.5435

0.6038 10.37 -0.6915

0.7030 10.80 -1.305

0.8028 11.81 -1.421

0.9006 13.45 -1.051

Acetone + water(T=308.15K)

0.5046 11.86 -1.538

0.6038 12.54 0.0635

37

0.7030 12.98 -0.4087

0.8028 13.81 0.4798

0.9006 16.41 -0.1791

Acetone + water(T=313.15K)

0.5046 13.94 -0.6440

0.6038 14.71 1.139

0.7030 15.38 0.4945

0.8028 16.02 0.8302

0.9006 19.26 0.8666

（1）x1 is the mole fraction of the solute, x2 is the mole fraction of water in binary mixtures.

Standard uncertainties (x1) = 0.01, (x2) = 0.0001, (p) =0.02, u (T) = 0.05 K,

( 2 ) ARD is the average relative deviation.

38

Table 7 Parameters of the modified Apelblat equation used to obtain solubility data

with different values of x2 for the binary mixtures

2x A B C AAD%

Water(2)+Methanol(3)

0.5099 572.62 -29270.2 -84.04 0.2380

0.6082 769.90 -38071.6 -113.5 0.4844

0.7044 871.91 -42817.0 -128.6 0.4749

0.8017 466.36 -24620.3 -68.16 0.7710

0.9049 422.56 -22696.5 -61.61 0.3197

Water(2)+ethanol(3)

0.8985 604.61 -30462.1 -88.96 0.3089

0.7984 198.04 -11952.1 -28.52 0.8435

0.6984 281.30 -15905.8 -40.82 0.6952

0.5990 65.330 -6145.30 -8.68 0.4673

0.4999 16.740 -3935.21 -1.470 0.2242

Water(2)+acetone(3)

0.5046 636.39 -32246.9 -93.51 0.6157

0.6038 714.14 -35368.7 -105.3 0.4954

0.7030 352.74 -19182.9 -51.45 0.5698

0.8028 256.36 -14739.8 -37.15 0.7895

0.9006 557.55 -28553.7 -81.90 0.5277

(1) The number in parenthesis, such as Water (2), (3) refers to solvents, such as methanol, ethanol,

and acetone.

(2) x2 is the mole fraction of water with the uncertainty of ± 0.0001

(3) AAD is the average absolute deviation.

39

Table 8 Thermodynamic functions related to the dissolution process of 5′-IMPNa2 in

binary mixtures containing organic solvents ( 2x ) at T = 303.15 K.

2x o

sso Hln∆ /(kJ·mol-1

) o

so Gln∆ /(kJ·mol-1

) o

so Sln∆ /J·K-1

·mol-1

) Hζ% TSζ%

Methanol + water

0.5099 31.5 10.5 69.5 59.96 40.04

0.6082 30.5 10.7 65.4 60.59 39.41

0.7044 31.8 10.8 69.4 60.20 39.80

0.8017 32.9 11.0 72.5 59.97 40.02

0.9049 33.4 11.1 73.6 59.96 40.04

Ethanol + water

0.4999 29.1 10.6 60.8 61.19 38.81

0.5990 27.8 11.0 55.6 62.27 37.72

0.6984 29.7 11.2 60.9 61.66 38.34

0.7984 29.6 11.4 59.9 61.94 38.06

0.8985 29.4 11.6 58.5 62.36 37.63

Acetone + water

0.5046 32.4 10.9 70.9 60.13 39.87

0.6038 28.6 11.3 57.1 62.29 37.71

0.7030 29.8 11.4 60.7 61.86 38.14

0.8028 28.9 11.5 57.3 62.47 37.53

0.9006 31.0 11.7 63.5 61.66 38.34

(1) X2 is water mole fraction with the uncertainty of ±0.0001.

(2) o

sso Hln∆ ,o

so Gln∆ , o

so Sln∆ is enthalpy, Gibbs energy, entropy of the solute, respectively.

Hζ% , TSζ% is the contribution of enthalpy and entropy to the Gibbs energy, respectively.

40

Graphical abstract:

41

Highlights:

1. Solubility of 5′-IMPNa2 in various solvents was studied for the first time.

2. The solubility could be ranked as follows: water > methanol > ethanol > acetone.

3. Modified Apelblat equation gave the best correlating results.

4. Mixing Gibbs free energies, enthalpies, and entropies were predicted.

5. Solubility data and equations can optimize the crystallization conditions.