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Solubility of trisodium citrate in water + methanol mixtures at various temperatures
Journal of Molecular Liquids 221 (2016) 166–170
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Solubility of trisodium citrate in water + methanol mixtures at various temperatures Jafar Soleymani a, Vahid Jouyban-Gharamaleki b, Karim Jouyban-Gharamaleki c, William E. Acree Jr. d, Ernst Kenndler e, Abolghasem Jouyban f,g,⁎ a
Drug Applied Research Center, Tabriz University of Medical Sciences, Tabriz, Iran School of Engineering Emerging Technologies, University of Tabriz, Tabriz 51664, Iran Liver and Gasterointestinal Diseases Research Center, Tabriz University of Medical Sciences, Tabriz 51664, Iran d Department of Chemistry, University of North Texas, Denton, TX 76203-5070, United States e Institute for Analytical Chemistry, Faculty of Chemistry, University of Vienna, Währingerstrasse 38, A 1090 Vienna, Austria f Pharmaceutical Analysis Research Center and Faculty of Pharmacy, Tabriz University of Medical Sciences, Tabriz, Iran g Kimia Idea Pardaz Azarbayjan (KIPA) Science Based Company, Tabriz University of Medical Sciences, Tabriz 51664, Iran b c
a r t i c l e
i n f o
Article history: Received 3 April 2016 Received in revised form 22 May 2016 Accepted 24 May 2016 Available online 26 May 2016 Keywords: Trisodium citrate Solubility prediction Binary mixtures Jouyban–Acree–van't Hoff model
a b s t r a c t A laser based automated setup was employed to experimentally determine the solubility of anhydrous trisodium citrate in binary mixtures of water and methanol under atmospheric pressure within the temperature range of 293.2 K to 313.3 K. The measured mole fraction solubility of citrate was around 5 × 10−2 in pure water and decreased by one order of magnitude to 5 × 10−3 at mole fraction of methanol of 0.3 in the binary mixture; within the investigated composition of the mixed solvent the solubility remained then at a low level and was about 10−4 at methanol mole fraction of 0.69. In all solvents the solubility increased with increasing temperature; the effect was more pronounced in the mixed solvents than in the pure aqueous solution. In order to extend the solubility data for practical applications, the experimental values were correlated to a semi-empirical combined model of Jouyban-Acree and van't Hoff, resulting in calculated solubilities which agreed with the experimental ones within a mean percentage deviation (MPD) of approximately 10.8%. © 2016 Elsevier B.V. All rights reserved.
1. Introduction In many areas of science and technology the solubility of a compound – the solute – in a solvent is of outstanding relevancy. Solubility is the amount of solute that can be dissolved in a given quantity of the solvent resulting in the saturated solution. The proper selection of the suitable single solvent or of the mixture of different solvents is a prerequisite for a rational planning of strategy and implementation of processes in which liquid phase reactions are involved, like solubilization, crystallization, purification, or enrichment of compounds by adsorption onto solid phases (solid phase extraction) or partition into liquid phases (liquid-liquid extraction), to mention only a few methodologies. Solubility plays also an important role in basic research of solution chemistry: it e.g. allows determining quantitatively in which solvent a dissolved uncharged solute is better stabilized in comparison to other solvents (or solvent mixtures). This quantitative measure, the standard Gibbs energy of transfer, or the so-called transfer activity coefficient, respectively, can directly be calculated from the ratio of the solubility of ⁎ Corresponding author at: Pharmaceutical Analysis Research Center and Faculty of Pharmacy, Tabriz University of Medical Sciences, Tabriz 51664, Iran. E-mail address: [email protected] (A. Jouyban).
http://dx.doi.org/10.1016/j.molliq.2016.05.077 0167-7322/© 2016 Elsevier B.V. All rights reserved.
the molecular solute in the different solvents (for details, see e.g. references [1–5] and the literature cited therein). For the dissolution of the solute either a single solvent or mixtures of solvents are used in practice. For analytical purposes, the organic co-solvents (also named modifiers) are applied to improve the resolution of the analytes by modifying properties decisive for the separation selectivity: partition coefficients and retention factors, electrophoretic mobilities, pKa values, etc. They are also applied to reduce the analysis time by decreasing the partition coefficients in liquid chromatography, either under isocratic conditions with solvent mixtures with constant composition, or by the application of gradients of the composition of the solvent mixture during the elution. It is important in this context that the modification of this composition might also influence the solubility of other compounds beside the analytes, namely matrix components of the sample, additives to the mobile phase like buffers, complexation agents, etc. Commonly the highest solubility of salts consisting of small ions – the compound of the present investigation – is observed in water as solvent; however, it is noteworthy to mention that a number of organic salts are more soluble in organic solvents, especially when these electrolytes consist of larger organic ions with low charge number. An interesting example in this context is a salt that is often used for physico-
J. Soleymani et al. / Journal of Molecular Liquids 221 (2016) 166–170
chemical investigations in solution chemistry: tetraphenylphosphonium tetraphenylborate. This salt consists of an equally sized, single charged cation and anion with low charge density. It is much more soluble in organic solvents like methanol (MeOH) or acetonitrile than in water [6]. However, the present paper deals with trisodium citrate consisting of a triple charged small organic anion that possesses an additional hydrophilic OH-group. This property leads to a high solubility in water, and we expected a significantly lower one in organic solvents, and as a consequence also in binary aqueous-methanolic mixtures. We selected trisodium citrate (1,2,3-propanetricarboxylic acid, 2hydroxy-, trisodium salt), also simply named sodium citrate, as solute, because it is an important compound used in many application areas [7], and we determined the change of its solubility as a function of the content of MeOH as organic co-solvent in binary mixtures. Trisodium citrate is produced on an industrial scale by the neutralization of citric acid with sodium hydroxide, and exists in crystallized form either as anhydrous, dihydrate or pentahydrate, respectively. From the wide application range (few examples are e.g. as retarder to modify desulfurization gypsum [8], as pre-cleaning agent used for monocrystalline silicon sheet [9], as cold-roll steel sheet passivator [10], as concrete modifier [11], or as environment-friendly degreasant for automobile coating [12], and in an enormous number of liquid mixtures used in technical processes) we point out here two main areas: the food and nutrition industry (here it is used as an additive), and the medical field. In food industry (the E number of the salt is E331) it is added for preservative and flavoring purposes. It is also used for pH regulation, because it can serve as a buffer over the pH range of 2 to 8 (it has pKa values of 3.13, 4.76 and 6.40 at 25 °C and zero ionic strength [13]), depending on the ratio of sodium hydroxide and the weak citric acid. In food industry it is also used as ingredient in many products, e.g. in beverages and soft drinks, in tinned fruit and vegetable products, milk powder, to name only a few. In the medical field sodium citrate is used to suppress the coagulation of blood samples in blood collection tubes and in blood banks. In addition to a number of various applications, it is also utilized e.g. for the treatment of infections of the urinary tract. Importantly it is major ingredient (together with glucose and sodium chloride) of the oral rehydration solution recommended by the WHO (the WHO ORS). As mentioned above, mixtures of sodium citrate with citric acid can serve as buffers, covering the pH range between about 2 to 8. When applied as a buffer or background electrolyte in separation methods like HPLC or capillary electrophoresis (CE), in many cases the solvent is not pure water, but consists of pure organic liquids or their mixtures, or mixtures of water with organic liquids. In both separation methods most common are aqueous binary mixtures of MeOH or acetonitrile, respectively, as organic co-solvents. This fact leads to the decisive question about the solubility of sodium citrate in such mixtures. In the present paper we investigated therefore the solubility of this compound in binary mixtures of water and MeOH. Solubilities were experimentally determined over a mole fraction of MeOH between zero (pure aqueous solution) and 0.31 for mixed solvents with 8 different compositions and at 5 temperatures between 293.2 K and 313.2 K. Based on the experimental data we applied a semi-empirical model of Jouyban-Acree and van't Hoff that enables calculating the solubilities at any temperature and solvent composition in the given range. This approach delivers a tool to predict data that certainly are useful in practice to select stable and reproducible liquid-phase systems. 2. Experimental section 2.1. Materials Trisodium citrate (trisodium 2-hydroxypropane-1,2,3-tricarboxylate, anhydrous, Na3C6H5O7, 258.06 g.mol−1, CAS Number 68-04-2), was purchased from BDH chemicals (England), methanol (purity of 0.999 m/m) from Scharlau Chemie (Spain). Double distilled water (from Shahid
167
Ghazi pharmaceutical company) was used throughout this study. All the materials were applied in the experiments as directly received from the supplier without any further purification. Table 1 provides full details of the chemicals. 2.2. Apparatus and methods Various methods have been applied for solubility determinations (reviewed by Jouyban and Fakhree [14]). In the present work a labmade system, whose principle and apparatus have been described in previously published papers [15], was used for determining the salt solubility in water + MeOH mixtures at temperatures of 293.2, 298.2, 303.2, 308.2 and 313.2 K. This setup is based on a synthetic method [16] and the disappearance upon dissolution of the solid solute added by a syringe is monitored by a laser beam. The experimental process could be briefly described as follows. Initially, the binary solvent mixtures were prepared by weighting the appropriate amounts of the solvents into an especially designed jacketed glass vessel whose temperature was controlled by a thermostat with the uncertainty of ± 0.1 K. Then, depending on the approximate amount of salt to be added, the appropriate syringe was selected and filled with trisodium citrate weighed using an electronic balance (Sartorius, Germany) with an uncertainty of 0.01 g. The syringe injects solid solute to the solvent mixture by receiving the command of the intelligent laser based sensor. To detect the saturation point, the setup checks repeatedly the number of particles suspended in solution to be higher than those particles detected in the solute-free mixture of solvent for about 2 h. Each experimental data is the average of at least three replicated measurements; the average value was used to calculate the mole fraction solubility; the resulting data were reproducible within 3.1%. 3. Calculation 3.1. Calculation of mole fraction The masses of the added amount of trisodium citrate were converted to mole fraction solubilities as following. First the according mole fraction of solvent i (1 for water and 2 for MeOH) in the absence of the solute, indicated by index 0 were computed as:
x0i
n ¼ X2 i
n i¼1 i
mi MW i
¼
m1 m2 þ MW 1 MW 2
ð1Þ
ni = mi/MWi, is the number of moles of solvent, i.e. m1 and m2 are the masses of the solvent 1 (water) and solvent 2 (MeOH) in the binary mixture, MW1 and MW2 are the molar masses of the respective solvents. The mole fraction solubilities at saturation (xm,T) of trisodium citrate in the various mixtures of the binary solvents at various temperatures were calculated by:
xm;T
mCitrate MW Citrate ¼ Xn mi mCitrate þ i¼2 MW MW Citrate i
ð2Þ
Table 1 List of the used materials. Material
Purity/in mass fraction
Company
Country
Trisodium citrate MeOH Water
0.999 0.999 conductivity b1.5 μS⋅ cm−1
BDH Scharlau Shahid Ghazi Pharmaceutical Co.
England Spain Iran
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in which mCitrate and MWCitrate are the mass and the molar mass, respectively, of trisodium citrate added to the solution. 3.2. Validating of models To express the accuracy of the model, the mean percentage deviation (MPD) between calculated and experimental solubilities was chosen; it is calculated as:
MPD ¼
Exp Cal 100 X xm;T −xm;T Exp N xm;T
! ð3Þ
Exp where N is the number of data points in each set; xCal m,T and xm,T refer to the calculated and the experimental solubilities.
4. Results and discussion 4.1. Experimental solubility data Table 2 lists the experimental mole fraction solubility of trisodium citrate in solutions with various mole fractions of binary mixtures of water (1) + MeOH (2) at 293.2, 298.2, 303.2, 308.2 and 313.2 K. To visualize the effect of the solvent on solubility, the data are depicted graphically in Fig. 1. Two tendencies were observed: first, the solubility in all tested solvent mixtures increased with increasing temperature, with the effect being larger in the MeOH-rich mixtures than in the water-rich ones. Secondly, with increasing MeOH concentration the solubilities decreased. In the double linear graph a steep decrease of xm,T is observed upon addition of MeOH up to its mole fraction of about 0.3. In this range, the solubility decreases from approximately 0.05 (in pure water) to around 0.005. Upon further increase of the MeOH concentration the solubility seemingly does not change as pronouncedly and remains at a low level. However, if we consider the data from Table 2 more closely, we can see that the relative decrease of the solubility with increasing MeOH concentration (i.e. the ratio of the decreasing solubilities related to a certain increase of the MeOH mole fraction) in fact is quite similar over the entire solvent composition range. This can be seen from the insert in Fig. 1, depicting the solubility vs. solvent composition in a semi-logarithmic plot. The resulting curves decrease steadily and moderately linear; the straight line depicts the linear correlation line at 293.2 K (R is 0.97). The curves of the logarithm of the solubility of the salt vs. the MeOH content have a roughly constant slope (they do not reach a plateau, with a small exception: in a short interval around the molar fractions of 0.5 the solubility changes only very slightly). We can derive from this curves that the citrate solubility decreases over the entire composition range approximately by one order of magnitude per increase of the MeOH mole fraction of 0.25. In analytical separation methods for practical reasons molar concentrations of the solutes are more common than mole fractions. This holds e.g. for the usage of citric acid/citrate buffers in mixed solvents as mobile
Fig. 1. Experimental solubilities, xm,T of trisodium citrate in dependence on the composition of the mixed solvents consisting of water (1) and MeOH (2) at different temperatures. Solvent composition is given in mole fraction of MeOH, x02. Insert: plot with xm,T in logarithmic scale. The straight line is the linear regression line for the data at T = 293.2 K; R = 0.97. Symbols for temperature: 293.2 K, □; 298.2 K, ○; 303.2 K, △; 308.2 K, ▽; 313.2 K, ◇.
phases in HPLC or background electrolytes in CE. From our data follows that the molar solubility of the citrate is relatively high in pure aqueous solution, in the present temperature range it is around 1.5 mol L−1. It is, in contrast, pronouncedly lower in the MeOH-rich fractions: it is e.g. only below the 10 mmol L−1 range at the mole fraction of MeOH of 0.6 to 0.7 with practical consequences e.g. of possible precipitation of the salt in the separation system (see below). 4.2. Prediction of solubility by models based on Jouyban-Acree and van't Hoff equations In order to extend the usage of the measured solubilities the Jouyban-Acree and van't Hoff equations were applied to correlate the experimental solubilities and to construct a trained model for the prediction of the solubilities to the entire temperature and solvent composition ranges under consideration. The van't Hoff equation is able to calculate the solubility data of the solute in each solvent system at various temperatures (xT) [17]:
ln xT ¼ A þ
B T
ð4Þ
T is the absolute temperature, A and B are the model constants calculated using a least square method. Based on Eq. (4), Fig. 2 shows the logarithm of the mole fraction solubilities versus the reciprocal absolute temperature in the investigated solvent systems at different temperatures. Indeed we can see the predicted linear dependence for all solvents.
Table 2 Experimental mole fraction solubility of trisodium citrate in binary mixtures of water (1) + MeOH (2) at various temperatures and atmospheric pressure (0.1 MPa). x0a 1
293.2 K
1.00 0.94 0.88 0.81 0.73 0.64 0.54 0.43 0.31
0.04326 0.04086 0.02278 0.01202 0.00354 0.00171 0.00144 0.00095 0.00005
SDb
298.2 K 0.00162 0.00058 0.00116 0.00050 0.00011 0.00007 0.00005 0.00003
b10−5
0.04406 0.04198 0.02467 0.01385 0.00482 0.00196 0.00160 0.00113 0.00007
SD
303.2 K 0.00196 0.00003 0.00070 0.00065 0.00025 0.00004 0.00005 0.00004
b10−5
0.04638 0.04309 0.02672 0.01504 0.00559 0.00221 0.00171 0.00124 0.00010
SD
308.2 K 0.00214 0.00053 0.00098 0.00069 0.00015 0.00006 0.00002 0.00005
b10−5
0.04951 0.04430 0.02918 0.01666 0.00686 0.00245 0.00189 0.00130 0.00013
The relative standard uncertainty for the solubilities is 3.1% or ur(x) = 0.031, the standard uncertainty for temperature is 0.1 K. a x01; mole fractions of water without solute. b SD: standard deviation.
SD
313.2 K 0.00249 0.00078 0.00061 0.00079 0.00007 0.00004 0.00004 0.00004
b10−5
0.05248 0.04569 0.03396 0.01817 0.00811 0.00279 0.00229 0.00136 0.00016
SD 0.00251 0.00102 0.00165 0.00041 0.00025 0.00012 0.00004 0.00005 b10−5
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The Jouyban-Acree model is, in addition, capable to train Eq. (6) using the minimum number of solubility data points at two temperatures (293.2 and 313.2 K) with 18 data points; then it can be utilized to predict the solubilities at other temperatures by: 1109 17088 þ x02 5:491− ln xm;T ¼ x01 0:562− T T " 2 # 18868x01 x02 x01 −x02 24942x01 x02 27077x01 x02 x01 −x02 þ − þ T T T ð8Þ
Fig. 2. Natural logarithm of the experimental mole fraction solubility, lnxm,T, of trisodium citrate in water (1) + MeOH (2) mixtures in dependence on the reciprocal absolute temperature, 1/T. Symbols for mole fraction of water, x01: 1.0, ■; 0.94,△; 0.88, ▽; 0.81, ○; 0.73, ▷; 0.64, □; 0.54,◁; 0.43, ◇; 0.31, ▼.
With the applied trained model, the obtained MPD value for this prediction is 10.2 ± 5.0% (N = 27) [25,26]. Regarding the possible formation of various polymorphs of the trisodium citrate, the MPD value is sufficiently low to be well suitable in practice; this is certainly an advantage of the Jouyban-Acree model. 5. Conclusions
The Jouyban-Acree model which is deduced from the Redlich-Kister equation [18] has been widely used to predict or back-calculate the experimental data points for any desired temperature and fraction of the binary mixture as following [19–21]: lnxm;T ¼ x01 ln x1;T þ x02 ln x2;T þ
" # 2 i x01 x02 X J i x01 −x02 T i¼0
ð5Þ
xm,T is the mole fraction solubility in the solvent mixtures at a given absolute temperature T, x01 and x02 are the mole fractions of solvents 1 and 2 in the solute-free solvent system. x1,T and x2,T are the mole fraction solubilities of trisodium citrate in the mono-solvents 1 and 2, respectively, and the Ji terms are the model constants (expressed in K unit), that reflect the solvent-solvent and solvent-solute interactions. The Ji terms are computed by a no-intercept least square analysis using a SPSS software by regressing ln xm,T − x01 ln x1,T − x02 ln x2,T against x0 x0 ðx0 −x0 Þ
x01 x02 x01 x02 ðx01 −x02 Þ T T ,
2
and 1 2 T1 2 [22]. By applying the van't Hoff equation, the Jouyban-Acree model could be rewritten as [23,24]: # " 0 0X i x x 2 0 B1 B2 lnxm;T ¼ x01 A1 þ þ x02 A2 þ þ 1 2 J i x1 −x02 T T T i¼0
ð6Þ
where A1, B1, A2, and B2 are the model constants. This model enables calculating the solute solubility in the mixed solvents at different temperatures. Unlike Eq. (5), Eq. (6) offers a model to predict the solubility without requiring those in the mono-solvents. By the aid of the experimental solubilities of trisodium citrate in the binary solvent mixtures, the generated data points were fitted to Eq. (6) and resulted in the model constants and J-terms given in Eq. (7) as following. 1052:7 17148 þ x02 5:741− ln xm;T ¼ x01 0:355− T T " 2 # 18922x01 x02 x01 −x02 24926x01 x02 27047x01 x02 x01 −x02 þ − þ T T T ð7Þ The result of the calculation gives a coefficient of determination of R2 = 0.999 between experimental and back-calculated solubilities in dependence on the solvent compositions. The MPD of the back calculated solubilities is 10.8 ± 5.1% (N = 45). Note that for the prediction of the solubility this approach only requires the knowledge of the temperature and the solvent fractions of interest.
The experimental solubility of anhydrous trisodium citrate in binary solvent mixtures of water and MeOH in the temperature range from 293.2 K to 313.2 K was determined by using a lab-made automated laser monitoring system. The composition of the solvents, expressed in mole fractions, ranged for water from 1 (pure water) to 0.31 (with the according range for MeOH from zero to 0.69). For a given solvent, the salt solubilities increased with increasing temperature, in water by about 20% from the lowest to the highest temperature, in the mixed solvents by up to 60%. Upon addition of MeOH up to a mole fraction of 0.3 a steep decrease of the solubility of the salt by one order of magnitude was observed from around 0.05 in pure water to about 0.005; in molar concentrations this is a decrease of more than 1 mol L−1, or as much as about 300 g salt per liter solution. With further increasing the MeOH content, the mole fraction solubility of the salt remained at the low level below 0.05, droping to the 10−4 range at 0.69 mol fraction of the organic co-solvent. Over the entire investigated composition range this is a decrease of 2 to 3 orders of magnitude, with the consequence that the salt might precipitate upon addition of a significant amount of MeOH to an aqueous solution. This behavior might be useful e.g. for purification of the salt by precipitation and re-crystallization. It might be, on the other hand, a disadvantage e.g. when sodium citrate is used in liquid chromatography as a buffering additive to the mobile phase at high initial water content [27], followed by applying a solvent gradient with increasing content of organic solvent [28]. The resulting low solubility could lead to precipitation of the salt in the separation system impeding its operational capability by clogging. This is the reason why in the papers dealing with HPLC (and CE) analyses only low citrate buffer concentrations (at the low millimolar level [29]) were used, when MeOH was co-solvent. Common gradients in HPLC apply MeOH contents from 20 to 80% v/v. For 20% v/v MeOH, its mole fraction is about 0.1, and the solubility of the salt is high in this mixed solvent: it is in the one molar range, causing none of the mentioned problems here. As the final mobile phase gradient composition is often chosen at MeOH contents of 80% v/v, the molar fraction of the organic co-solvent is thus around 0.6. This is the composition where the solubility vs. solvent composition curves show a sharp change to the low solubility level (see Fig. 2 and the discussion given above). At this point the solubility is below the 10 millimolar range, in the critical range of solubility. To be on the safe side with the operational conditions in order to avoid interferences of the analysis due a too low solubility of citrate, mobile phases are used that consist of a number of additional co-solvents, mostly acetonitrile [30], tetrahydrofuran [31], buffers (like phosphate), or other additives [28,29,32]. However, literature shows that citrate buffer is not a common buffer used in these analytical separation methods (in CE it is in fact only rarely used) [33,34], although it might have some advantages: negligible absorbance in the UV range suitable
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for the most common UV absorbance detectors, or a wide buffering range. In order to obtain data of solubilities at any temperature and solvent composition within the investigated range, an approach is used in the present paper that allows calculating these solubilities. This was carried out by applying the combined form of Jouyban-Acree and van't Hoff equations to the measured data. This approach led to a correlation between measured and back-calculated solubilities with an MPD of 10.8%, a value that is well suited for practical demands. References [1] K. Sarmini, E. Kenndler, Ionization constants of weak acids and bases in organic solvents, J. Biochem. Biophys. Methods 38 (1999) 123–137. [2] E. Kenndler, A critical overview of non-aqueous capillary electrophoresis. Part I: mobility and separation selectivity, J. Chromatogr. A 1335 (2014) 16–30. [3] A. Jouyban, E. Kenndler, Capillary electrophoresis with organic solvents in pharmaceutical analysis: a systematic guide through the background, Curr. Anal. Chem. 10 (2014) 248–266. [4] E. Kenndler, Organic solvents in CE, Electrophoresis 30 (2009) S101–S111. [5] S.P. Porras, E. Kenndler, Capillary zone electrophoresis in non-aqueous solutions: pH of the background electrolyte, J. Chromatogr. A 1037 (2004) 455–465. [6] J. Muzikár, V.D. Goor, B. Gas, E. Kenndler, Electrophoretic mobilities of large organic ions in propylene carbonate, N,N′-dimethylformamide, N,N′-dimethylacetamide, acetonitrile and methanol: Determination by capillary electrophoresis in non-aqueous solvents, Electrophoresis 23 (2002) 375–382. [7] D. Sackett, Citric Acid: Occurrence, Biochemistry, Applications and Processing, Nova Science Publishers, Inc., New York, 2014. [8] F. Niu, H.A. Guo, X.W. Feng, Y. Meng, Modification of retarder on desulfurization gypsum wall material, Adv. Mater. Res. 741 (2013) 45–48. [9] Q. Wu, Pre-Cleaning Agent Used for Monocrystalline Silicon Sheet Before Texturing and Its Usage, Zhejiang Topoint Photovoltaic Co., Ltd., China, 2012. [10] Z. Yang, A Cold-Roll Steel Sheet Passivator, and Its Usage, Anhui Feida Advanced Material Technology Co., Ltd., Peop. Rep. China, 2015. [11] Q. Yang, Concrete Modifier and Its Usage, Tongji University, China, 2012. [12] Y. Ma, T. Pang, P. Cheng, F. Lang, X. Huang, Z. Shi, Y. Chen, R. He, EnvironmentFriendly Degreasant for Automobile Coating and Its Usage, Wuhan Iron and Steel Group Corp, China, 2016. [13] A. Albert, E.P. Serjeant, The Determination of Ionization Constants: A Laboratory Manual, Chapman and Hall, London, 1984. [14] A. Jouyban, M.A.A. Fakhree, in: W.E. Acree Jr. (Ed.), Toxicity and drug testing, InTech Publisher, New York, 2012. [15] V. Jouyban-Gharamaleki, K. Jouyban-Gharamaleki, A. Shayanfar, M. Khoubnasabjafari, A. Jouyban, An automated system for determination of drug solubility based on laser monitoring technique, J. Lab. Autom. 20 (2015) 3–9.
[16] R.W. Hankinson, D. Thompson, Equilibria and solubility data for methanol–water-1nitropropane ternary liquid system, J. Chem. Eng. Data 10 (1965) 18–19. [17] D.J.W. Grant, M. Mehdizadeh, A.H.L. Chow, J.E. Fairbrother, Non-linear van't Hoff solubility-temperature plots and their pharmaceutical interpretation, Int. J. Pharm. 18 (1984) 25–38. [18] W.E. Acree Jr., Mathematical representation of thermodynamic properties. Part II. Derivation of the combined nearly ideal binary solvent (NIBS)/Redlich-Kister mathematical representation from a two-body and three-body interactional mixing model, Thermochim. Acta 198 (1992) 71–79. [19] A. Jouyban-Gharamaleki, W.E. Acree Jr., Comparison of models for describing multiple peaks in solubility profiles, Int. J. Pharm. 167 (1998) 177–182. [20] A. Jouyban, Review of the cosolvency models for predicting solubility of drugs in water-cosolvent mixtures, J. Pharm. Pharm. Sci. 11 (2008) 32–58. [21] A. Jouyban, Handbook of Solubility Data for Pharmaceuticals, CRC Press, Boca Raton, FL, 2010. [22] A. Jouyban-Gharamaleki, J. Hanaee, A novel method for improvement of predictability of the CNIBS/R-K equation, Int. J. Pharm. 154 (1997) 245–247. [23] A. Jouyban, A. Shayanfar, W.E. Acree Jr., Solubility prediction of polycyclic aromatic hydrocarbons in non-aqueous solvent mixtures, Fluid Phase Equilib. 293 (2010) 47–58. [24] A. Shayanfar, S.H. Eghrary, F. Sardari, W.E. Acree Jr., A. Jouyban, Solubility of anthracene and phenanthrene in ethanol +2,2,4-trimethylpentane mixtures at different temperatures, J. Chem. Eng. Data 56 (2011) 2290–2294. [25] A. Beerbower, P.L. Wu, A. Martin, Expanded solubility parameter approach 1. Naphthalene and benzoic acid in individual solvents, J. Pharm. Sci. 73 (1984) 179–188. [26] A. Reillo, M. Cordoba, B. Escalera, E. Selles, M. Cordoba Jr., Prediction of sulfamethiazole solubility in dioxane-water mixtures, Pharmazie 50 (1995) 472–475. [27] M. Stadlober, M. Sager, K.J. Irgolic, Identification and quantification of selenium compounds in sodium selenite supplemented feeds by HPLC-ICP-MS, Bodenkultur 52 (2001) 233–241. [28] L.J. Burroughs, P.D.A. Williams, Single HPLC method for complete separation of unmodified and reduced iso-α acids, Proc. Congr. Eur. Brew. Conv. 27 (1999) 283–290. [29] A. Kumar, A.K. Malik, D.K. Tewary, B. Singh, Gradient HPLC of antibiotics in urine, ground water, chicken muscle, hospital wastewater, and pharmaceutical samples using C18 and RP-amide columns, J. Sep. Sci. 31 (2008) 294–300. [30] G.H. Yousefi, S.M. Foroutan, A. Zarghi, A.R. Shafaati, Quantification of polyethylene glycol esters of methotrexate and determination of their partition coefficients by validated HPLC methods, Iran J. Pharm. Res. 8 (2009) 27–31. [31] G.J. Lawson, J. Chakraborty, M.C. Dumasia, E.M. Baylis, Methylprednisolone hemisuccinate and metabolites in urine from patients receiving high-dose corticosteroid therapy, Ther. Drug Monit. 14 (1992) 20–26. [32] H. Tomankova, J. Sabartova, Determination of potential degradation products of piroxicam by HPTLC densiometry and HPLC, Chromatographia 28 (1989) 197–202. [33] C.E. Lin, C.C. Chang, W.C. Lin, Migration behavior and separation of sulfonamides in capillary zone electrophoresis. III. Citrate buffer as a background electrolyte, J. Chromatogr. A 768 (1997) 105–112. [34] L. Chang-Hai, S. Qing-Yan, S. Chun-Quan, Z. Hai, C. Jun, Z. Liang, C. Yi-Feng, Capillary electrophoresis for analysis of optical impurity of L-pramipexole, AJSMMU 30 (2009) 1058–1061.
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Journal of Molecular Liquids journal homepage: www.elsevier.com/locate/molliq
Solubility of trisodium citrate in water + methanol mixtures at various temperatures Jafar Soleymani a, Vahid Jouyban-Gharamaleki b, Karim Jouyban-Gharamaleki c, William E. Acree Jr. d, Ernst Kenndler e, Abolghasem Jouyban f,g,⁎ a
Drug Applied Research Center, Tabriz University of Medical Sciences, Tabriz, Iran School of Engineering Emerging Technologies, University of Tabriz, Tabriz 51664, Iran Liver and Gasterointestinal Diseases Research Center, Tabriz University of Medical Sciences, Tabriz 51664, Iran d Department of Chemistry, University of North Texas, Denton, TX 76203-5070, United States e Institute for Analytical Chemistry, Faculty of Chemistry, University of Vienna, Währingerstrasse 38, A 1090 Vienna, Austria f Pharmaceutical Analysis Research Center and Faculty of Pharmacy, Tabriz University of Medical Sciences, Tabriz, Iran g Kimia Idea Pardaz Azarbayjan (KIPA) Science Based Company, Tabriz University of Medical Sciences, Tabriz 51664, Iran b c
a r t i c l e
i n f o
Article history: Received 3 April 2016 Received in revised form 22 May 2016 Accepted 24 May 2016 Available online 26 May 2016 Keywords: Trisodium citrate Solubility prediction Binary mixtures Jouyban–Acree–van't Hoff model
a b s t r a c t A laser based automated setup was employed to experimentally determine the solubility of anhydrous trisodium citrate in binary mixtures of water and methanol under atmospheric pressure within the temperature range of 293.2 K to 313.3 K. The measured mole fraction solubility of citrate was around 5 × 10−2 in pure water and decreased by one order of magnitude to 5 × 10−3 at mole fraction of methanol of 0.3 in the binary mixture; within the investigated composition of the mixed solvent the solubility remained then at a low level and was about 10−4 at methanol mole fraction of 0.69. In all solvents the solubility increased with increasing temperature; the effect was more pronounced in the mixed solvents than in the pure aqueous solution. In order to extend the solubility data for practical applications, the experimental values were correlated to a semi-empirical combined model of Jouyban-Acree and van't Hoff, resulting in calculated solubilities which agreed with the experimental ones within a mean percentage deviation (MPD) of approximately 10.8%. © 2016 Elsevier B.V. All rights reserved.
1. Introduction In many areas of science and technology the solubility of a compound – the solute – in a solvent is of outstanding relevancy. Solubility is the amount of solute that can be dissolved in a given quantity of the solvent resulting in the saturated solution. The proper selection of the suitable single solvent or of the mixture of different solvents is a prerequisite for a rational planning of strategy and implementation of processes in which liquid phase reactions are involved, like solubilization, crystallization, purification, or enrichment of compounds by adsorption onto solid phases (solid phase extraction) or partition into liquid phases (liquid-liquid extraction), to mention only a few methodologies. Solubility plays also an important role in basic research of solution chemistry: it e.g. allows determining quantitatively in which solvent a dissolved uncharged solute is better stabilized in comparison to other solvents (or solvent mixtures). This quantitative measure, the standard Gibbs energy of transfer, or the so-called transfer activity coefficient, respectively, can directly be calculated from the ratio of the solubility of ⁎ Corresponding author at: Pharmaceutical Analysis Research Center and Faculty of Pharmacy, Tabriz University of Medical Sciences, Tabriz 51664, Iran. E-mail address: [email protected] (A. Jouyban).
http://dx.doi.org/10.1016/j.molliq.2016.05.077 0167-7322/© 2016 Elsevier B.V. All rights reserved.
the molecular solute in the different solvents (for details, see e.g. references [1–5] and the literature cited therein). For the dissolution of the solute either a single solvent or mixtures of solvents are used in practice. For analytical purposes, the organic co-solvents (also named modifiers) are applied to improve the resolution of the analytes by modifying properties decisive for the separation selectivity: partition coefficients and retention factors, electrophoretic mobilities, pKa values, etc. They are also applied to reduce the analysis time by decreasing the partition coefficients in liquid chromatography, either under isocratic conditions with solvent mixtures with constant composition, or by the application of gradients of the composition of the solvent mixture during the elution. It is important in this context that the modification of this composition might also influence the solubility of other compounds beside the analytes, namely matrix components of the sample, additives to the mobile phase like buffers, complexation agents, etc. Commonly the highest solubility of salts consisting of small ions – the compound of the present investigation – is observed in water as solvent; however, it is noteworthy to mention that a number of organic salts are more soluble in organic solvents, especially when these electrolytes consist of larger organic ions with low charge number. An interesting example in this context is a salt that is often used for physico-
J. Soleymani et al. / Journal of Molecular Liquids 221 (2016) 166–170
chemical investigations in solution chemistry: tetraphenylphosphonium tetraphenylborate. This salt consists of an equally sized, single charged cation and anion with low charge density. It is much more soluble in organic solvents like methanol (MeOH) or acetonitrile than in water [6]. However, the present paper deals with trisodium citrate consisting of a triple charged small organic anion that possesses an additional hydrophilic OH-group. This property leads to a high solubility in water, and we expected a significantly lower one in organic solvents, and as a consequence also in binary aqueous-methanolic mixtures. We selected trisodium citrate (1,2,3-propanetricarboxylic acid, 2hydroxy-, trisodium salt), also simply named sodium citrate, as solute, because it is an important compound used in many application areas [7], and we determined the change of its solubility as a function of the content of MeOH as organic co-solvent in binary mixtures. Trisodium citrate is produced on an industrial scale by the neutralization of citric acid with sodium hydroxide, and exists in crystallized form either as anhydrous, dihydrate or pentahydrate, respectively. From the wide application range (few examples are e.g. as retarder to modify desulfurization gypsum [8], as pre-cleaning agent used for monocrystalline silicon sheet [9], as cold-roll steel sheet passivator [10], as concrete modifier [11], or as environment-friendly degreasant for automobile coating [12], and in an enormous number of liquid mixtures used in technical processes) we point out here two main areas: the food and nutrition industry (here it is used as an additive), and the medical field. In food industry (the E number of the salt is E331) it is added for preservative and flavoring purposes. It is also used for pH regulation, because it can serve as a buffer over the pH range of 2 to 8 (it has pKa values of 3.13, 4.76 and 6.40 at 25 °C and zero ionic strength [13]), depending on the ratio of sodium hydroxide and the weak citric acid. In food industry it is also used as ingredient in many products, e.g. in beverages and soft drinks, in tinned fruit and vegetable products, milk powder, to name only a few. In the medical field sodium citrate is used to suppress the coagulation of blood samples in blood collection tubes and in blood banks. In addition to a number of various applications, it is also utilized e.g. for the treatment of infections of the urinary tract. Importantly it is major ingredient (together with glucose and sodium chloride) of the oral rehydration solution recommended by the WHO (the WHO ORS). As mentioned above, mixtures of sodium citrate with citric acid can serve as buffers, covering the pH range between about 2 to 8. When applied as a buffer or background electrolyte in separation methods like HPLC or capillary electrophoresis (CE), in many cases the solvent is not pure water, but consists of pure organic liquids or their mixtures, or mixtures of water with organic liquids. In both separation methods most common are aqueous binary mixtures of MeOH or acetonitrile, respectively, as organic co-solvents. This fact leads to the decisive question about the solubility of sodium citrate in such mixtures. In the present paper we investigated therefore the solubility of this compound in binary mixtures of water and MeOH. Solubilities were experimentally determined over a mole fraction of MeOH between zero (pure aqueous solution) and 0.31 for mixed solvents with 8 different compositions and at 5 temperatures between 293.2 K and 313.2 K. Based on the experimental data we applied a semi-empirical model of Jouyban-Acree and van't Hoff that enables calculating the solubilities at any temperature and solvent composition in the given range. This approach delivers a tool to predict data that certainly are useful in practice to select stable and reproducible liquid-phase systems. 2. Experimental section 2.1. Materials Trisodium citrate (trisodium 2-hydroxypropane-1,2,3-tricarboxylate, anhydrous, Na3C6H5O7, 258.06 g.mol−1, CAS Number 68-04-2), was purchased from BDH chemicals (England), methanol (purity of 0.999 m/m) from Scharlau Chemie (Spain). Double distilled water (from Shahid
167
Ghazi pharmaceutical company) was used throughout this study. All the materials were applied in the experiments as directly received from the supplier without any further purification. Table 1 provides full details of the chemicals. 2.2. Apparatus and methods Various methods have been applied for solubility determinations (reviewed by Jouyban and Fakhree [14]). In the present work a labmade system, whose principle and apparatus have been described in previously published papers [15], was used for determining the salt solubility in water + MeOH mixtures at temperatures of 293.2, 298.2, 303.2, 308.2 and 313.2 K. This setup is based on a synthetic method [16] and the disappearance upon dissolution of the solid solute added by a syringe is monitored by a laser beam. The experimental process could be briefly described as follows. Initially, the binary solvent mixtures were prepared by weighting the appropriate amounts of the solvents into an especially designed jacketed glass vessel whose temperature was controlled by a thermostat with the uncertainty of ± 0.1 K. Then, depending on the approximate amount of salt to be added, the appropriate syringe was selected and filled with trisodium citrate weighed using an electronic balance (Sartorius, Germany) with an uncertainty of 0.01 g. The syringe injects solid solute to the solvent mixture by receiving the command of the intelligent laser based sensor. To detect the saturation point, the setup checks repeatedly the number of particles suspended in solution to be higher than those particles detected in the solute-free mixture of solvent for about 2 h. Each experimental data is the average of at least three replicated measurements; the average value was used to calculate the mole fraction solubility; the resulting data were reproducible within 3.1%. 3. Calculation 3.1. Calculation of mole fraction The masses of the added amount of trisodium citrate were converted to mole fraction solubilities as following. First the according mole fraction of solvent i (1 for water and 2 for MeOH) in the absence of the solute, indicated by index 0 were computed as:
x0i
n ¼ X2 i
n i¼1 i
mi MW i
¼
m1 m2 þ MW 1 MW 2
ð1Þ
ni = mi/MWi, is the number of moles of solvent, i.e. m1 and m2 are the masses of the solvent 1 (water) and solvent 2 (MeOH) in the binary mixture, MW1 and MW2 are the molar masses of the respective solvents. The mole fraction solubilities at saturation (xm,T) of trisodium citrate in the various mixtures of the binary solvents at various temperatures were calculated by:
xm;T
mCitrate MW Citrate ¼ Xn mi mCitrate þ i¼2 MW MW Citrate i
ð2Þ
Table 1 List of the used materials. Material
Purity/in mass fraction
Company
Country
Trisodium citrate MeOH Water
0.999 0.999 conductivity b1.5 μS⋅ cm−1
BDH Scharlau Shahid Ghazi Pharmaceutical Co.
England Spain Iran
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in which mCitrate and MWCitrate are the mass and the molar mass, respectively, of trisodium citrate added to the solution. 3.2. Validating of models To express the accuracy of the model, the mean percentage deviation (MPD) between calculated and experimental solubilities was chosen; it is calculated as:
MPD ¼
Exp Cal 100 X xm;T −xm;T Exp N xm;T
! ð3Þ
Exp where N is the number of data points in each set; xCal m,T and xm,T refer to the calculated and the experimental solubilities.
4. Results and discussion 4.1. Experimental solubility data Table 2 lists the experimental mole fraction solubility of trisodium citrate in solutions with various mole fractions of binary mixtures of water (1) + MeOH (2) at 293.2, 298.2, 303.2, 308.2 and 313.2 K. To visualize the effect of the solvent on solubility, the data are depicted graphically in Fig. 1. Two tendencies were observed: first, the solubility in all tested solvent mixtures increased with increasing temperature, with the effect being larger in the MeOH-rich mixtures than in the water-rich ones. Secondly, with increasing MeOH concentration the solubilities decreased. In the double linear graph a steep decrease of xm,T is observed upon addition of MeOH up to its mole fraction of about 0.3. In this range, the solubility decreases from approximately 0.05 (in pure water) to around 0.005. Upon further increase of the MeOH concentration the solubility seemingly does not change as pronouncedly and remains at a low level. However, if we consider the data from Table 2 more closely, we can see that the relative decrease of the solubility with increasing MeOH concentration (i.e. the ratio of the decreasing solubilities related to a certain increase of the MeOH mole fraction) in fact is quite similar over the entire solvent composition range. This can be seen from the insert in Fig. 1, depicting the solubility vs. solvent composition in a semi-logarithmic plot. The resulting curves decrease steadily and moderately linear; the straight line depicts the linear correlation line at 293.2 K (R is 0.97). The curves of the logarithm of the solubility of the salt vs. the MeOH content have a roughly constant slope (they do not reach a plateau, with a small exception: in a short interval around the molar fractions of 0.5 the solubility changes only very slightly). We can derive from this curves that the citrate solubility decreases over the entire composition range approximately by one order of magnitude per increase of the MeOH mole fraction of 0.25. In analytical separation methods for practical reasons molar concentrations of the solutes are more common than mole fractions. This holds e.g. for the usage of citric acid/citrate buffers in mixed solvents as mobile
Fig. 1. Experimental solubilities, xm,T of trisodium citrate in dependence on the composition of the mixed solvents consisting of water (1) and MeOH (2) at different temperatures. Solvent composition is given in mole fraction of MeOH, x02. Insert: plot with xm,T in logarithmic scale. The straight line is the linear regression line for the data at T = 293.2 K; R = 0.97. Symbols for temperature: 293.2 K, □; 298.2 K, ○; 303.2 K, △; 308.2 K, ▽; 313.2 K, ◇.
phases in HPLC or background electrolytes in CE. From our data follows that the molar solubility of the citrate is relatively high in pure aqueous solution, in the present temperature range it is around 1.5 mol L−1. It is, in contrast, pronouncedly lower in the MeOH-rich fractions: it is e.g. only below the 10 mmol L−1 range at the mole fraction of MeOH of 0.6 to 0.7 with practical consequences e.g. of possible precipitation of the salt in the separation system (see below). 4.2. Prediction of solubility by models based on Jouyban-Acree and van't Hoff equations In order to extend the usage of the measured solubilities the Jouyban-Acree and van't Hoff equations were applied to correlate the experimental solubilities and to construct a trained model for the prediction of the solubilities to the entire temperature and solvent composition ranges under consideration. The van't Hoff equation is able to calculate the solubility data of the solute in each solvent system at various temperatures (xT) [17]:
ln xT ¼ A þ
B T
ð4Þ
T is the absolute temperature, A and B are the model constants calculated using a least square method. Based on Eq. (4), Fig. 2 shows the logarithm of the mole fraction solubilities versus the reciprocal absolute temperature in the investigated solvent systems at different temperatures. Indeed we can see the predicted linear dependence for all solvents.
Table 2 Experimental mole fraction solubility of trisodium citrate in binary mixtures of water (1) + MeOH (2) at various temperatures and atmospheric pressure (0.1 MPa). x0a 1
293.2 K
1.00 0.94 0.88 0.81 0.73 0.64 0.54 0.43 0.31
0.04326 0.04086 0.02278 0.01202 0.00354 0.00171 0.00144 0.00095 0.00005
SDb
298.2 K 0.00162 0.00058 0.00116 0.00050 0.00011 0.00007 0.00005 0.00003
b10−5
0.04406 0.04198 0.02467 0.01385 0.00482 0.00196 0.00160 0.00113 0.00007
SD
303.2 K 0.00196 0.00003 0.00070 0.00065 0.00025 0.00004 0.00005 0.00004
b10−5
0.04638 0.04309 0.02672 0.01504 0.00559 0.00221 0.00171 0.00124 0.00010
SD
308.2 K 0.00214 0.00053 0.00098 0.00069 0.00015 0.00006 0.00002 0.00005
b10−5
0.04951 0.04430 0.02918 0.01666 0.00686 0.00245 0.00189 0.00130 0.00013
The relative standard uncertainty for the solubilities is 3.1% or ur(x) = 0.031, the standard uncertainty for temperature is 0.1 K. a x01; mole fractions of water without solute. b SD: standard deviation.
SD
313.2 K 0.00249 0.00078 0.00061 0.00079 0.00007 0.00004 0.00004 0.00004
b10−5
0.05248 0.04569 0.03396 0.01817 0.00811 0.00279 0.00229 0.00136 0.00016
SD 0.00251 0.00102 0.00165 0.00041 0.00025 0.00012 0.00004 0.00005 b10−5
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The Jouyban-Acree model is, in addition, capable to train Eq. (6) using the minimum number of solubility data points at two temperatures (293.2 and 313.2 K) with 18 data points; then it can be utilized to predict the solubilities at other temperatures by: 1109 17088 þ x02 5:491− ln xm;T ¼ x01 0:562− T T " 2 # 18868x01 x02 x01 −x02 24942x01 x02 27077x01 x02 x01 −x02 þ − þ T T T ð8Þ
Fig. 2. Natural logarithm of the experimental mole fraction solubility, lnxm,T, of trisodium citrate in water (1) + MeOH (2) mixtures in dependence on the reciprocal absolute temperature, 1/T. Symbols for mole fraction of water, x01: 1.0, ■; 0.94,△; 0.88, ▽; 0.81, ○; 0.73, ▷; 0.64, □; 0.54,◁; 0.43, ◇; 0.31, ▼.
With the applied trained model, the obtained MPD value for this prediction is 10.2 ± 5.0% (N = 27) [25,26]. Regarding the possible formation of various polymorphs of the trisodium citrate, the MPD value is sufficiently low to be well suitable in practice; this is certainly an advantage of the Jouyban-Acree model. 5. Conclusions
The Jouyban-Acree model which is deduced from the Redlich-Kister equation [18] has been widely used to predict or back-calculate the experimental data points for any desired temperature and fraction of the binary mixture as following [19–21]: lnxm;T ¼ x01 ln x1;T þ x02 ln x2;T þ
" # 2 i x01 x02 X J i x01 −x02 T i¼0
ð5Þ
xm,T is the mole fraction solubility in the solvent mixtures at a given absolute temperature T, x01 and x02 are the mole fractions of solvents 1 and 2 in the solute-free solvent system. x1,T and x2,T are the mole fraction solubilities of trisodium citrate in the mono-solvents 1 and 2, respectively, and the Ji terms are the model constants (expressed in K unit), that reflect the solvent-solvent and solvent-solute interactions. The Ji terms are computed by a no-intercept least square analysis using a SPSS software by regressing ln xm,T − x01 ln x1,T − x02 ln x2,T against x0 x0 ðx0 −x0 Þ
x01 x02 x01 x02 ðx01 −x02 Þ T T ,
2
and 1 2 T1 2 [22]. By applying the van't Hoff equation, the Jouyban-Acree model could be rewritten as [23,24]: # " 0 0X i x x 2 0 B1 B2 lnxm;T ¼ x01 A1 þ þ x02 A2 þ þ 1 2 J i x1 −x02 T T T i¼0
ð6Þ
where A1, B1, A2, and B2 are the model constants. This model enables calculating the solute solubility in the mixed solvents at different temperatures. Unlike Eq. (5), Eq. (6) offers a model to predict the solubility without requiring those in the mono-solvents. By the aid of the experimental solubilities of trisodium citrate in the binary solvent mixtures, the generated data points were fitted to Eq. (6) and resulted in the model constants and J-terms given in Eq. (7) as following. 1052:7 17148 þ x02 5:741− ln xm;T ¼ x01 0:355− T T " 2 # 18922x01 x02 x01 −x02 24926x01 x02 27047x01 x02 x01 −x02 þ − þ T T T ð7Þ The result of the calculation gives a coefficient of determination of R2 = 0.999 between experimental and back-calculated solubilities in dependence on the solvent compositions. The MPD of the back calculated solubilities is 10.8 ± 5.1% (N = 45). Note that for the prediction of the solubility this approach only requires the knowledge of the temperature and the solvent fractions of interest.
The experimental solubility of anhydrous trisodium citrate in binary solvent mixtures of water and MeOH in the temperature range from 293.2 K to 313.2 K was determined by using a lab-made automated laser monitoring system. The composition of the solvents, expressed in mole fractions, ranged for water from 1 (pure water) to 0.31 (with the according range for MeOH from zero to 0.69). For a given solvent, the salt solubilities increased with increasing temperature, in water by about 20% from the lowest to the highest temperature, in the mixed solvents by up to 60%. Upon addition of MeOH up to a mole fraction of 0.3 a steep decrease of the solubility of the salt by one order of magnitude was observed from around 0.05 in pure water to about 0.005; in molar concentrations this is a decrease of more than 1 mol L−1, or as much as about 300 g salt per liter solution. With further increasing the MeOH content, the mole fraction solubility of the salt remained at the low level below 0.05, droping to the 10−4 range at 0.69 mol fraction of the organic co-solvent. Over the entire investigated composition range this is a decrease of 2 to 3 orders of magnitude, with the consequence that the salt might precipitate upon addition of a significant amount of MeOH to an aqueous solution. This behavior might be useful e.g. for purification of the salt by precipitation and re-crystallization. It might be, on the other hand, a disadvantage e.g. when sodium citrate is used in liquid chromatography as a buffering additive to the mobile phase at high initial water content [27], followed by applying a solvent gradient with increasing content of organic solvent [28]. The resulting low solubility could lead to precipitation of the salt in the separation system impeding its operational capability by clogging. This is the reason why in the papers dealing with HPLC (and CE) analyses only low citrate buffer concentrations (at the low millimolar level [29]) were used, when MeOH was co-solvent. Common gradients in HPLC apply MeOH contents from 20 to 80% v/v. For 20% v/v MeOH, its mole fraction is about 0.1, and the solubility of the salt is high in this mixed solvent: it is in the one molar range, causing none of the mentioned problems here. As the final mobile phase gradient composition is often chosen at MeOH contents of 80% v/v, the molar fraction of the organic co-solvent is thus around 0.6. This is the composition where the solubility vs. solvent composition curves show a sharp change to the low solubility level (see Fig. 2 and the discussion given above). At this point the solubility is below the 10 millimolar range, in the critical range of solubility. To be on the safe side with the operational conditions in order to avoid interferences of the analysis due a too low solubility of citrate, mobile phases are used that consist of a number of additional co-solvents, mostly acetonitrile [30], tetrahydrofuran [31], buffers (like phosphate), or other additives [28,29,32]. However, literature shows that citrate buffer is not a common buffer used in these analytical separation methods (in CE it is in fact only rarely used) [33,34], although it might have some advantages: negligible absorbance in the UV range suitable
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