determination of solubility of campna in water + (ethanol, methanol, and acetone) within...

Determination of Solubility of cAMPNa in Water + (Ethanol,Methanol, and Acetone) within 293.15−313.15 KPengpeng Yang, Qingshi Wen, Jinglan Wu, Wei Zhuang, Yuehui Zhang, and Hanjie Ying*

National Engineering Technique Research Center for Biotechnology, State Key Laboratory of Materials-Oriented ChemicalEngineering, College of Biotechnology and Pharmaceutical Engineering, Nanjing Tech University, Nanjing 210009, PR China

*S Supporting Information

ABSTRACT: The solubility of adenosine 3′,5′-cyclic monophosphate sodium (cAMPNa) in mixed solvents (water+ethanol,water+methanol, and water+acetone) was measured within 293.15−313.15 K under atmospheric pressure. The (CNIBS)/Redlich−Kister model and the modified Apelblat equation were respectively applied to correlate the solubility data to evaluatethe effect of the compositional ratio of the organic solvent and the temperature on the solubility in binary solvents, andsatisfactory simulation results were obtained. The solubility of cAMPNa was maximal in pure water and markedly diminished atall evaluated temperatures as the mole fraction of the organic solvent in the aqueous mixture increased. The thermodynamicfunctions for cAMPNa dissolution in the three solvent mixtures were obtained from the solubility data using the van’t Hoff andGibbs equations, and the dissolution behavior was discussed. Dissolution of cAMPNa was endothermic and nonspontaneous inall cases, and the enthalpy was the major contributing force to the Gibbs energy.

1. INTRODUCTION

Adenosine 3′,5′-cyclic monophosphate (cAMP, CAS RegistryNo. 60-92-4, molecular mass: 329.2) plays an important role inchanges in the nervous system,1 cardiovascular system,2 andclinically treated diseases such as myocardial infarction,myocarditis, and angina.3 Commercial cAMP is primarilyproduced by chemical synthesis, and several industrialpurification and separation processes are utilized to isolatethe drug from its side-products. Previously, a novel method forcAMP biofermentation and the corresponding separation andpurification methods were reported by the present researchgroup,4 achieving dramatic cost reduction compared with thatof traditional chemical synthesis.5 Regardless of the approachutilized, the crystallization process, which is one of the last andmost important processes, is vital for obtaining a high-qualityproduct.The solubility behavior of a drug is instrumental in the design

and optimization of a reasonable crystallization process,6 andthus, it is also important for administration of the drug tohumans.7 Consequently, it is necessary to determine thesolubility of a compound as a function of temperature andsolvent composition. However, to date, data related to thesolubility of cAMP or cAMPNa (Figure 1) has scarcely been

reported. Some solubility data are available for similarnucleotide salts such as inosine-5′-monophosphate disodium,8

cytidine 5′-monophosphate disodium,9 and cytidine 5′-diphosphocholine disodium10 in water and different binarymixtures. The acid form of cAMP is slightly soluble in water,whereas cAMPNa is effortlessly dissolved in water but is poorlysoluble in organic solvents such as alcohols. cAMP exists mainlyin the salt form in the fermentation broth and thus is usuallyconverted to the cAMPNa form in industrial purification.The balance and dynamic methods are the most commonly

used methods for the determination of solubility. Theseapproaches may be further classified as UV spectrophoto-metric11 and gravimetric methods12 based on differences in thedetection technique. The gravimetric approach is applicable toa wide range of functions, whereas the UV spectrophotometricapproach offers higher accuracy and convenience. Lasermonitoring13 is generally applied to the dynamic method, andis more sensitive for the identification of phase equilibrium andis suitable for systems that rapidly achieve equilibrium. Becausea relatively long time is required for cAMPNa to attain thesolid−liquid equilibrium, the balance method based onspectrophotometry (referring to the isothermal method utilizedherein) may be adequate for determining the solubility ofcAMPNa in organic−water mixtures. The solubility of theaforementioned nucleotides has been measured by adoptingthis approach. Acree et al.14 proposed the Redlich−Kistermodel for describing how the experimental isothermal solubilityof a crystalline solute dissolved in a binary solvent mixturevaries with the composition of the binary solvent. The modifiedApelblat equation based on the solid−liquid phase equilibrium

Received: March 8, 2014Revised: May 25, 2014Accepted: May 29, 2014Published: May 29, 2014

Figure 1. Molecular structure of cAMPNa.

Article

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© 2014 American Chemical Society 10803 dx.doi.org/10.1021/ie501000k | Ind. Eng. Chem. Res. 2014, 53, 10803−10809

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theory15 was a widely used semiempirical model and providesvery accurate mathematical descriptions for solid−liquidequilibrium data.16 These models have been applied tobiological systems to correlate the solubility data with thesolvent ratio and temperature in pure solvents and variousbinary solvents, leading to satisfactory results. The model hasalso been applied to pioglitazone hydrochloride,17 losartanpotassium,18 lovastatin,19 cloxacillin sodium,20 and theaforementioned nucleotides. Besides, for the ternary systemscontaining electrolyte/amino acid/water solutions, one ad-vanced model, the electrolyte perturbed-chain statisticalassociation theory (ePC-SAFT), may be applied to modelthermodynamic properties.21 Christoph Held modeled andpredicted the amino-acid solubility in the ternary salt/aminoacid/water system with an overall deviation from experimentaldata of 9.3%.22

Herein, the solubility of cAMPNa was determined in variousbinary mixed solvents (methanol + water, ethanol + water, andacetone + water) within 293.15−313.15 K under atmosphericpressure by using an isothermal method. The (CNIBS)/Redlich-Kister model and the modified Apelblat equation wereselected for correlating the solubility of cAMPNa with thesolvent composition and temperature. The magnitude of thethermodynamic parameters for the solution process wasdetermined from the solubility data, in addition to the analysisof the enthalpy−entropy compensation to discuss thedissolution mechanism and behavior. This study aims toprovide theoretical solubility data as a platform for manipulat-ing the cAMPNa crystallization process.

2. EXPERIMENTAL SECTION

2.1. Materials. The cAMPNa used for determining thesolubility in the present experiment contains two crystallizationwater molecules (cAMPNa·2H2O) with a purity of greater than0.99 in mass fraction. This sample was prepared byrecrystallization from an ethanol + water system. cAMPNa·2H2O (CCDC No. 976996) crystallizes with trigonalsymmetry; the infrared spectra (Nexus 670, Thermo ElectronCorporation, USA) and the X-ray powder diffraction(apparatus type ARL X′TRA, Thermo Electron Corporation,USA) data based on our recent study are respectively shown inFigures 2 and 3.Analytical-grade methanol, ethanol, and acetone (mass

fractions higher than 0.995) were obtained from TianjinKewei Chemical Reagent Co., Ltd., China, and used withoutfurther purification.2.2. Solubility Measurements. The solubility of cAMPNa

in water + methanol, water + ethanol, and water + acetonemixtures within 293.15−313.15 K was measured by theisothermal method described elsewhere.9 In brief, an excessamount of cAMPNa·2H2O and aqueous mixtures with suitablecompositions were first added to a conical flask and agitatedwith a magnetic stirrer for 4 h to reach equilibrium. Theexperiments were carried out in a constant-temperature waterbath (type DC-2030, Shanghai Sunny Hengping ScientificInstrument Co. Ltd., China) with the desired temperaturemaintained within ±0.05 K. Stirring was discontinued after 10h, and the solution was allowed to stand for a further 3 h. Theupper portion of the solution was taken and filtered (underisothermal conditions), and the concentration was determinedby measuring the UV absorption at 260 nm (UV-2000,UNICO, USA) after appropriate dilution.

All of the solubility experiments were performed in triplicate(three independently obtained results). The mean values werepresented and used in the calculations. The initial mole fractioncomposition of the binary solvent mixtures is defined by eq 1,and the mole fraction solubility of cAMPNa in water + ethanol,water + methanol, and water + acetone within 283.15−323.15K was calculated by using eq 2:

=+

xm M

m M m M( / )

( / ) ( / )22 2

1 1 2 2 (1)

=+ +

xm M

m M m M m M( / )

( / ) ( / ) ( / )33 3

1 1 2 2 3 3 (2)

where x3 represents the solubility (solid−liquid equilibria) ofcAMPNa in mole fraction; x2 represents the solute-free molefraction of antisolvent in the binary liquid solvents; m1, m2, andm3, respectively, represent the mass of water (taking the crystalwater of cAMPNa·2H2O into account), organic solvent(ethanol, methanol, and acetone), and solute (without regardto crystal water) in different solvent systems; M1, M2, and M3,respectively, represent the molecular weight of water, relevantsolvents, and solute.

Figure 2. Infrared spectrum of cAMPNa·2H2O.

Figure 3. X-ray power diffraction pattern of cAMPNa·2H2O.

Industrial & Engineering Chemistry Research Article

dx.doi.org/10.1021/ie501000k | Ind. Eng. Chem. Res. 2014, 53, 10803−1080910804

3. RESULTS AND DISCUSSIONThe solubility data for cAMPNa in water + ethanol, water +methanol, and water + acetone binary mixed solvents within293.15−313.15 K are summarized in the SupportingInformation. In these binary solvents, the mole fraction oforganic solvent (ethanol, methanol, and acetone) ranges from 0to 0.5. On the basis of the solubility data, within thetemperature range under consideration, the solubility ofcAMPNa in the three types of mixed solvent increases withincrease in the temperature and is reduced as the mole fractionof the organic solvent increases.3.1. Solubility Data and (CNIBS)/Redlich−Kister Model

Correlation. To correlate the experimental solubility with thesolvent composition, Acree et al.14,23,24 suggested the combinednearly ideal binary solvent (CNIBS)/Redlich−Kister model asa possible mathematical representation for describing how theexperimental isothermal solubility of a crystalline solutedissolved in a solvent mixture varies with solvent composition.For binary solvent systems, the model can be rearranged to eq3:

= + + + +x B B x B x B x B xln( ) ( ) ( ) ( )3 0 1 2 2 22

3 23

4 24

(3)

The solubility data for the binary system of water + (ethanol,methanol, or acetone) are correlated by eq 3, and the calculatedsolubilities and values of parameters B0, B1, B2, B3, and B4 forthe model are listed in the Supporting Information along withthe relevant root-mean-square deviation (RMSD). Additionally,the solubility data calculated using the (CNIBS)/Redlich−Kister model are in good agreement with the experimentalvalues (Figure 4).The relative deviation (RD) and RMSD are defined as25,11

=−x xx

RD 3 3cal

3 (4)

∑= −=

⎡⎣⎢⎢

⎤⎦⎥⎥N

x xRMSD1

( )i

N

13 3

cal 21/2

(5)

where N is the number of experimental points, x3cal represents

the calculated solubility values, and x3 indicates theexperimental solubility values. The RD and RMSD werecalculated from eq 4 and eq 5 to assess the accuracy of thecorrelation models.The solubility of cAMPNa was reduced in different

concentrations of the three binary mixtures relative to that inpure water as the proportion of the organic solvent increased.The solubility of cAMPNa in the three solvent systemsfollowed the general order: water + acetone < water + ethanol< water + methanol < pure water. cAMPNa belongs to the classof polyhydroxyl ionic compounds and exists in aqueoussolution in the form of cAMP− and Na+. The oxygen atomsof the −PO groups, −OH groups in the ribose moiety, andN atom of the adenine base in the molecular structure ofcAMP− act, in solution, as Lewis bases to establish a largenumber of hydrogen bonds with electron-acceptor functionalgroups in the aqueous solution (such as the hydrogen atoms inH2O). cAMP− can also act as an electron acceptor owing to theH atoms of the ribose moiety and adenine unit to establishhydrogen bonds with the oxygen atoms of H2O. Thisinteraction results in a hydration layer around cAMP− similarto that of Na+ in aqueous solutions, so that cAMPNa exhibitsfavorable solubility in water-rich solvent mixtures. Furthermore,

the high dielectric constant (εr = 80 at 293.15 K for water)26

leads to shielding of the ions (Na+ and cAMP−) from eachother, further contributing to the good solubility of the salt inwater.In contrast, in the organic-rich mixed solvents, because of the

smaller dielectric constant, εr, of the organic material (acetone:21.4, ethanol: 24.6, and methanol: 32.7 at 293.15 K) relative towater, the hydrophobicity of the molecule tends to play animportant role in spite of the existing hydrogen bonds betweencAMP− and the solvent molecules (including the organics andwater), thus increasing the repulsion between cAMP− and thesolvent with a consequent decline in the solubility of the solute,

Figure 4. Mole fraction solubility of cAMPNa (x3) in solvent mixture(a) ethanol (2) + water (1), (b) methanol (2) + water (1), and (c)acetone (2) + water (1) at various temperatures: (☆) T = 293.15 K;(▽) T = 298.15 K; (△) T = 303.15 K; (○) T = 308.15 K; (□) T =313.15 K. Points represent the experimental data. Curves arecalculated according to eq 3 using the (CNIBS)/Redlich−Kistermodel.



which is in accord with the general rule “like dissolves like.”Because of the lack of literature data, no other directcomparison of the cAMPNa solubility values in the threeaqueous systems is available. In addition, the three organicsolvents can be used as antisolvents for dilution crystallizationof cAMPNa; ethanol may be recommended as the better choicebased on safety considerations.3.2. Modified Apelblat Equation Correlation. On the

basis of the solid−liquid phase equilibrium theory, the modifiedApelblat equation can also be used to correlate the solubility ofcAMPNa with the temperature as follows:13,14,16,27

= + +xBT

C Tln( ) A ln( )3 (6)

where A, B, and C are empirical constants. The values of A andB represent the variation in the solution activity coefficient, andthe C value reflects the effect of temperature on the fusionenthalpy. The experimental solubility values were fitted to eq 6by the least-squares method,28 and the results are shown inFigure 5. The values of the three parameters a, b, and c arelisted in the Supporting Information along with the RMSD.The solubility of cAMPNa in the three aqueous mixtures

positively correlates with the temperature; that is, the solubilityincreases as the temperature rises, particularly in the lowerorganic solvent concentration region where the change is morepronounced compared to the slight variation above 0.4 molefraction of the three organic solvents (Figure 5). Thus, toenhance the yield, the use of a cooling process is not effective atlater stages of the crystallization process if these three organicsolvents are used as industrial antisolvents in the dilutioncrystallization of cAMPNa.In addition, the calculated solubility data correlated by the

modified Apelblat equation are in good agreement with theexperimental solubility data. Thus, the equation model can beused to calculate the solubility of cAMPNa in the three solventmixtures to generate satisfactorily accurate results.3.3. Thermodynamic Functions of the Dissolution

Process. For nonideal solutions, the solubility can be expressedby the van’t Hoff equation modified by the appropriate activitycoefficient as follows, eq 7:29,16

γ = −Δ

−⎛⎝⎜

⎞⎠⎟x

HR T T

ln( )1 1

3f

f (7)

where x1 is the mole fraction of the solute in the solution, T isthe solution temperature (K), Tf is the fusion temperature(melting point) of the solute (K), ΔHf is the molar enthalpy offusion of the solute (J mol−1), R is the gas enthalpy constant(8.314 J mol−1 K−1), and γ is the activity coefficient.In such cases, the enthalpy and entropy of mixing must be

taken into account by replacing ΔHf with ΔHsol (where solindicates “of dissolution”) and ΔSf by ΔSsol.30,17 In addition,over a limited temperature interval (293.15−313.15 K) thechange in the heat capacity of the solution can be assumed tobe constant;31,18 hence, the derived values of ΔHsol would bevalid for the harmonic mean experimental temperature (Thm =n/Σ(1/Ti) = 302.99 K).32 Thus, the van’t Hoff equation can betransformed and expressed in the form of eq 8, and the changein the molar enthalpy of dissolution, ΔHsol, can be calculatedusing eq 9 (in which the slope indicated in the equation refersto that of the ln x1 vs (1/T − 1/Thm) plot). Thus

Δ = − ∂∂

= − ∂∂ −

⎛⎝⎜

⎞⎠⎟

⎛⎝⎜

⎞⎠⎟H R

xT

Rx

T Tln( )

(1/ )ln( )

(1/ 1/ )p p

solhm

1/2

(8)

Figure 5. Solubility of cAMPNa (x3) in solvent mixture (a) ethanol(2) + water (1), (b) methanol (2) + water (1), (c) acetone (2) +water (1) correlated with different temperatures: (□) x2 = 0.1; (○) x2= 0.2; (△) x2 = 0.3; (▽) x2 = 0.3; (★) x2 = 0.4; (●) x2 = 0.5. Thepoints represent the experimental data. Curves are calculatedaccording to eq 6 using the Modified Apelblat equation.



−Δ = ×H R slopesol (9)

Figure 6 shows some typical plots for cAMPNa in the threesolvent mixtures at 302.99 K as examples. In this study, the

relationship between ln(x3) and (1/T − 1/Thm) is linear for thethree solvent systems under all conditions under consideration.The molar Gibbs energy change for the dissolution process,ΔGsol, can be calculated from the intercept of a regression plotof ln(x3) against (1/T − 1/Thm), which can be expressed as eq10:33

Δ = ×G RT interceptsol hm (10)

The molar entropy change of dissolution can be obtainedfrom the Gibbs equation and expressed as eq 11:34

Δ =Δ − Δ

SH G

Tsolsol sol

hm (11)

To quantitatively estimate the relative contribution of thedissolution of cAMPNa in the aqueous mixtures to the Gibbsenergy, the enthalpy (%H) and entropy (%TS) calculated fromeqs 12 and 13 are introduced.35,36

ξ =|Δ |

|Δ | + | Δ |H

H T S% 100H

sol

sol sol (12)

ξ =| Δ |

|Δ | + | Δ |T S

H T S% 100TS

sol

sol sol (13)

The relative contributions of the enthalpy %ξH and entropy%ξTS to the Gibbs free energy of dissolution were obtainedfrom eqs 12 and 13. ΔHsol, ΔGsol, ΔSsol, %ξH, and %ξTS aresummarized in Table 1. The propagation of uncertainties in thethermodynamic quantities was calculated according to theliterature method.37,24 The enthalpy change of dissolution waspositive in all cases; therefore, the process is alwaysendothermic, which is consistent with the fact that thesolubility of cAMPNa in water increases with increasingtemperature. The entropy of dissolution in all three types ofsolvent mixtures, and in neat water, was positive, whichindicates that the entropy is the overall driving force for the

dissolution process for all the mixtures studied. The ΔHsolvalues mainly increase at lower concentrations of the organicsolvent and then decline as the concentration of organic solventincreases. ΔSsol follows the same tendency.From Table 1, it is apparent that in all the mixtures, the main

contributor to the molar Gibbs energy of dissolution ofcAMPNa is the enthalpy, particularly for pure water and 0.5-mole fraction of acetone aqueous solution and ethanol aqueoussolution. In addition, ΔGsol exhibits an increasing trend in theorganic solvent mole fraction range 0−0.5 for the dissolution ofcAMPNa in these aqueous mixtures, which indicates that thedissolution process is not facile, which agrees with the empiricalfacts. Furthermore, the positive ΔGsol value obtained in all thesystems indicates that the dissolution process was notspontaneous; that is, in practical execution of this process,the system absorbs more than energy from the surroundings.Analysis of the enthalpy−entropy compensation (Figure 7)

allows deduction of the mechanism of cAMPNa dissolution. Achange in the slope of the plot indicates a change in thedominant mechanism.38,25 For the water + acetone and water +methanol mixtures, the slope was positive in the organic solventmole fraction range 0−0.3, indicating that the dominantmechanism was enthalpy-controlled. However, the negativeslope in the mole fraction range 0.3−0.5, demonstrates achange to an entropy-controlled process. A similar tendencywas observed for the water+ethanol mixture, where the drivingforce changed from enthalpic to entropic at 0.2 mole fraction ofethanol.

4. CONCLUSIONThe solubility of cAMPNa in water + ethanol, water +methanol, and water + acetone mixtures was analyzed by theisothermal method within 293.15−313.15 K. The experimentaldata were effectively correlated by the (CNIBS)/Redlich−Kister model and the modified Apelblat equation. It can be

Figure 6. Temperature dependence of solubility of cAMPNa (x3) indifferent solvent mixtures: (△) water + methanol (2), x2 = 0.4; (○)water + acetone (2), x2 = 0.3; (□) water + ethanol (2), x2 = 0.4.

Table 1. Thermodynamic Functions of Dissolution ofcAMPNa (3) in Water(1) + Ethanol(2), Water (1) +Methanol (2), and Water (1) + Acetone(2) Mixtures at Thm= 302.99 K

x2 molefraction

ΔHsol (kJ/mol)

ΔGsol (kJ/mol)

ΔSsol (J/(mol·K)) %ζH %ζTS

Pure Water0.000 13.97 7.73 20.60 69.12 30.88

Water + Ethanol0.100 29.09 8.48 67.98 58.53 41.470.200 37.73 9.52 93.05 57.22 42.780.300 31.62 11.09 67.75 60.63 39.370.400 27.62 12.59 49.59 64.76 35.240.500 21.43 14.35 23.36 75.16 24.84

Water + Methanol0.100 24.12 8.48 51.58 60.67 39.330.200 36.10 9.57 87.52 57.64 42.360.300 48.26 11.08 122.64 56.48 43.520.400 40.98 12.20 94.93 58.74 41.260.500 38.72 13.07 84.59 60.16 39.84

Water + Acetone0.100 27.15 8.51 61.50 59.29 40.710.200 35.10 9.84 83.32 58.15 41.850.300 40.76 11.62 96.10 58.32 41.680.400 40.07 13.60 87.32 60.22 39.780.500 29.70 15.42 47.10 67.53 32.47



concluded that the dissolution of cAMPNa in the three solventmixtures, and in neat water, depends strongly on the solventcomposition and temperature. The solubility of cAMPNa isfavorable in pure water and declines as the proportion of theorganic solvent increases. Ethanol may be a better choice as anantisolvent for dilution crystallization of cAMPNa thanmethanol or acetone; cooling does not exert much effect onthe enhancement of the yield coefficient at later stages of theprocess. The thermodynamic functions (enthalpy, entropy, andmolar Gibbs free energy) of dissolution were calculated fordifferent solvent mixtures based on the solubility data. Thesolubility proved to be consistently endothermic and notspontaneous for the dissolution of cAMPNa in the threesolvent mixtures throughout the concentration range studied,and the enthalpy was the major contributing force to the Gibbsenergy. For the three binary systems, cAMPNa dissolution wasinitially enthalpy-controlled followed by a change in themechanism of dissolution to entropy-controlled in the organicsolvent mole fraction range 0−0.5, based on the analysis of theenthalpy−entropy compensation. The experimental solubilityand correlation equation provided in this work can be used forreference and research of the crystallization of cAMPNa.

■ ASSOCIATED CONTENT

*S Supporting InformationDetails concerning the chemicals, experimental solubility ofcAMPNa, and optimized parameters for CNIBS/Redlich−Kister model and modified Apelblat equation. This material isavailable free of charge via the Internet at http://pubs.acs.org.

■ AUTHOR INFORMATION

Corresponding Author*Tel.: +86 25 86990001. Fax: +86 25 58133398. E-mail:[email protected].

NotesThe authors declare no competing financial interest.

■ ACKNOWLEDGMENTS

This work was supported by the National Basic ResearchProgram of China (No. 2011CBA00806), Natural ScienceFoundation of Jiangsu Grants (No. BK20130929, BK2011031),Jiangsu Postdoctoral Science Foundation (1301038B), NationalOutstanding 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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Figure 7. ΔHsol vs ΔGsol enthalpy−entropy compensation plot fordissolution of cAMPNa in three solvent mixtures at 302.99 K: (■)water + ethanol; (●) water + methanol; (▲) water + acetone.



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determination of solubility of campna in water + (ethanol, methanol, and acetone) within...

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