experimental measurement and modelling of solubility of inosine-5′-monophosphate disodium in pure...
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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
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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: [email protected]
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: [email protected]
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).
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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.