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Solubility and Solution Thermodynamics of SomeSulfonamides in 1-Propanol + Water Mixtures
Daniel R. Delgado • Fleming Martınez
Received: 28 October 2013 / Accepted: 23 January 2014 / Published online: 10 May 2014� Springer Science+Business Media New York 2014
Abstract The solubilities of sulfadiazine (SD), sulfamerazine (SMR) and sulfamethazine
(SMT) in some 1-propanol ? water co-solvent mixtures were measured at five temperatures
from 293.15 to 313.15 K over the polarity range provided by the aqueous solvent mixtures.
The mole fraction solubility of all these sulfonamides was maximal in the 0.80 mass fraction
of 1-propanol solvent mixture (dsolv = 28.3 MPa1/2) and minimal in water (d = 47.8 MPa1/2)
at all temperatures studied. The apparent thermodynamic functions Gibbs energy, enthalpy,
and entropy of solution were obtained from these solubility data by using the van’t Hoff
and Gibbs equations. Apparent thermodynamic quantities of mixing were also calculated
by using the ideal solubilities reported in the literature. Nonlinear enthalpy–entropy rela-
tionships were observed for these drugs in the plots of enthalpy versus Gibbs energy of
mixing. The plot of DmixH� versus DmixG� shows different trends according to the slopes
obtained when the mixture compositions change. Accordingly, the mechanism for the
solution process of SD and SMT in water-rich mixtures is enthalpy driven, whereas it is
entropy driven for SMR. In a different way, in 1-propanol-rich mixtures the mechanism is
enthalpy driven for SD and SMR and entropy driven for SMT. Ultimately, in almost all of
the intermediate compositions, the mechanism is enthalpy driven. Nevertheless, the
molecular events involved in the solution processes remain unclear.
Keywords Sulfonamides � 1-Propanol ? water mixtures � Solubility � Solution
thermodynamics � Activity coefficient
1 Introduction
The behavior of drugs in different co-solvent mixtures as function of temperature and
composition is studied for the main purposes of substances purification, design of liquid
D. R. Delgado � F. Martınez (&)Grupo de Investigaciones Farmaceutico-Fisicoquımicas, Departamento de Farmacia, UniversidadNacional de Colombia, A.A. 14490 Bogota D.C., Colombiae-mail: fmartinezr@unal.edu.co
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J Solution Chem (2014) 43:836–852DOI 10.1007/s10953-014-0169-0
medicine formulations, and understanding the mechanisms involved in the physical and
chemical stabilization of dissolved pharmaceuticals [1, 2]. Therefore, the equilibrium
solubility of drugs is a very important physicochemical property to be considered during
the design of pharmaceutical products because it affects the drug efficacy by influencing
several biopharmaceutical and pharmacokinetic properties [3]. On the other hand, the
temperature dependence of the solubility allows a thermodynamic analysis in order to gain
insight into the molecular mechanisms involved in the drug dissolution processes [4].
Sulfadiazine (SD, molar mass 250.28 g�mol-1, Fig. 1), sulfamerazine (SMR, molar mass
264.31 g�mol-1, Fig. 1) and sulfamethazine (SMT, molar mass 278.33 g�mol-1, Fig. 1) are
sulfonamide drugs that are used as effective antimicrobial agents for the prevention and cure
of several kinds of bacterial infections in human and veterinary therapies [5].
Although SD, SMR and SMT have been widely used in therapeutics, solubility data for
these drugs in co-solvent mixtures is scarce in the literature [6]. Even though some theoretical
and semiempirical models have been developed to predict drug solubilities, the availability of
experimental data is still fundamental information for pharmaceutical scientists [7]. Because
the solubility of sulfonamides in neat water is so low [8, 9], co-solvent ? water mixtures have
been evaluated in order to increase the solubility of some of these compounds [10, 11]. These
studies have also been carried out to understand the molecular mechanisms involved in the
drug dissolution processes. In this regard, a complete physicochemical study has recently been
presented to contribute to the understanding of the mechanisms involved in the transfer of some
sulfonamides between aqueous and organic systems [12].
On the other hand, although 1-propanol is not widely used as co-solvent for design of
liquid medicines, it has been used as solvent in the pharmaceutical industry for resins and
cellulose esters [8]. This co-solvent is a hydrogen-donor and hydrogen-acceptor due to its
hydroxyl group, and thus it is miscible with water in all proportions [13].
The main goal of this work is thus to extend the database of experimental solubilities for
SD, SMR and SMT, and also to evaluate the effect of the co-solvent composition on
solubility and solution thermodynamics of these drugs in binary mixtures formed with
1-propanol and water, based on the van’t Hoff method, including the respective contri-
butions from mixing of these solvent compounds toward the solution processes, as has
been done with other sulfonamides in different co-solvent systems [14–17]. Therefore, this
thermodynamic study is very similar to the ones reported previously about the solubility of
SD, SMR and SMT in ethanol ? water mixtures [18, 19].
2 Experimental
2.1 Materials
Sulfadiazine (SD, component 3, CAS [68-35-9], 4-amino-N-pyrimidin-2-yl-benzene-
sulfonamide, with purity greater than 0.990 in mass fraction), sulfamerazine (SMR,
NH2
SNH
N
N
OO
R1
R2
Fig. 1 Molecular structure ofthe sulfonamides studied.Sulfadiazine: R1 and R2 = H;sulfamerazine: R1 = H,R2 = CH3; sulfamethazine: R1
and R2 = CH3
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component 3, CAS [127-79-7], 4-amino-N-(4-methylpyrimidin-2-yl)benzenesulfonamide,
with purity greater than 0.990 in mass fraction) and sulfamethazine (SMT, component 3,
CAS [57-68-1], 4-amino-N-(4,6-dimethylpyrimidin-2-yl)benzenesulfonamide, with purity
greater than 0.990 in mass fraction) from Sigma Chemical Company, 1-propanol A.R. from
Merck (component 1, with purity greater than 0.995 in mass fraction), and distilled water
(component 2) with conductivity \2 lS�cm-1, were used in this research. A molecular
sieve (Merck, number 3, pore size 0.3 nm) and Durapore� filters (0.45 lm, Millipore
Corp.) were also used. The source and purities of the compounds (expressed in mass
fractions) used in this work are summarized in Table 1.
2.2 Solvent Mixture Preparations
All 1-propanol ? water solvent mixtures were prepared by mass, using an Ohaus Pioneer
TM PA214 analytical balance with sensitivity ±0.1 mg, in quantities of 50.00 g. In order
to cover the entire composition range, the mass fractions of propanol, w1, of the nine binary
mixtures prepared was varied by 0.10 from 0.10 to 0.90.
2.3 Solubility Determinations
The procedures used in this research are similar to those employed previously in the study
of these sulfonamides in ethanol ? water mixtures [18, 19]. Briefly, an excess of SD, SMR
or SMT was added to approximately 10 g of each co-solvent mixture or neat solvents in
stoppered dark glass flasks. The flasks with the solid–liquid mixture were placed in an
ultrasonic bath (Elma� E60H Elmasonic, USA) for 15 min and later they were placed in
thermostatic mechanical shakers (Julabo SW23, Germany) kept at 303.15, 308.15, or
313.15 (±0.05) K, and also in re-circulating thermostatic baths (Neslab RTE 10 Digital
One Thermo Electron Company, USA) kept at 293.15 or 298.15 (±0.05) K, for at least
4 days to reach equilibrium. This equilibrium time was established by measuring the drug
concentrations in neat water at 293.15 K until they became constant. It is expected that the
equilibration times in the co-solvent mixtures will be smaller due to the greater drug
solubilities in these systems. After this time the supernatant solutions were filtered under
isothermal conditions (Millipore Corp. Swinnex�-13, USA) to ensure that they were free
of particulate matter before sampling.
Drug concentrations were determined after appropriate alcoholic dilution by measuring
the UV light absorbance at 268 nm for all drugs (UV/VIS BioMate 3 Thermo Electron
Company spectrophotometer, USA) and interpolation from previously constructed UV
spectrophotometric calibration curves. All the solubility experiments were run at least in
triplicate. In order to transform mole fractions to molar concentrations (mol�dm-3), the
density of the saturated solutions were determined by using a digital density meter (DMA
45 Anton Paar, Austria) connected to the same re-circulating thermostatic baths according
to procedures described in the literature [20].
3 Results and Discussion
Before proposing the possible intermolecular interactions present in the saturated solutions
of SD, SMR or SMT, it is important to keep in mind that these drugs in solution act mainly
as Lewis bases (due to their –NH2, –SO2–, and =N– groups) and as Lewis acids (due to
838 J Solution Chem (2014) 43:836–852
123
their –NH2 and[N–H groups) in order to establish hydrogen bonds with the –OH groups
in the solvents [18, 19, 21].
3.1 Experimental and Ideal Solubility
Tables 2 and 3 list the experimental solubilities (expressed in mole fraction and molarity,
respectively) of SD, SMR and SMT in 1-propanol ? water mixtures over the temperature
range studied, 293.15–313.15 K. In almost all cases the variation coefficients of the solubility
for all of the drugs were smaller than 2.0 %. Comparison of sulfonamide solubilities in neat
water with respect to values reported in the literature has been made previously [18, 19].
Some differences are found between data reported by Zhang et al. and data reported in
Table 2 for the mole fraction solubility of SD in neat 1-propanol at four of the temperatures
studied here, our values being almost three times lower than those reported previously (i.e.
1.243 9 10-4 at 298.15 K, 1.642 9 10-4 at 303.15 K, 2.085 9 10-4 at 308.15 K, and
2.446 9 10-4 at 313.15 K, respectively) [22]. Similarly, big differences are found between
data reported by Zhang et al. in another paper and data reported in Table 2 for the mole
fraction solubility of SMT in neat 1-propanol at all the temperatures studied here, with our
values being larger by almost two orders of magnitude than those reported previously (i.e.
2.746 9 10-6 at 293.15 K, 4.966 9 10-6 at 298.15 K, 8.314 9 10-6 at 303.15 K,
1.276 9 10-5 at 308.15 K, and 1.846 9 10-5 at 313.15 K, respectively) [23]. These
discrepancies could be due to different polymorphic solid states, different saturation times,
or different analytical techniques, among others, as has been described in the literature [6].
Finally, to the best of our knowledge, no solubility values have been published for SMR in
neat 1-propanol or for any these drugs in 1-propanol ? water mixtures, and therefore no
other comparison is possible.
The solubility increases with temperature in all cases indicating that the dissolution
process is endothermic. The highest mole fraction solubilities of SD, SMR and SMT were
obtained in the mixture with 0.80 in mass fraction of 1-propanol at T = 313.15 K, whereas
the lowest values are found in water at 293.15 K (Table 2). Nevertheless, if the molarity
(mol�dm-3) concentration scale is considered then the maximum solubilities are obtained
in the mixture 0.60 in mass fraction of 1-propanol for SD and 0.70 in mass fraction of
1-propanol for SMR and SMT (Table 3).
Table 2 also shows the ideal solubilities in mole fraction of the solutes (xid3 ) reported in
the literature [18, 19]. This table shows that the ideal solubilities of SD, SMR and SMT are
greater than the experimental solubilities obtained at all the temperatures studied. This
result could be explained as a consequence of the larger solvent–solvent and/or solute–
Table 1 Source and purities of the compounds used in this work
Compound CAS Formula Molar mass/g�mol-1
Source Purity in massfraction
Sulfadiazine 68-35-9 C10H10N4O2S 250.28 Sigma Chemical Co. 0.990
Sulfamerazine 127-79-7 C11H12N4O2S 264.31 Sigma Chemical Co. 0.990
Sulfamethazine 57-68-1 C12H14N4O2S 278.33 Sigma Chemical Co. 0.990
1-propanol 71-23-8 C3H8O 60.10 Merck 0.998
Water 7732-18-5 H2O 18.02 Obtained bydistillation
Conductivity\2 lS�cm-1
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Table 2 Experimental solubility of sulfadiazine, sulfamerazine and sulfamethazine in 1-propanol ? watermixtures expressed in mole fraction (104 x3) including ideal solubility at several temperatures (±0.05 K)
w1a x1
a 104 x3
T = 293.15 K T = 298.15 K T = 303.15 K T = 308.15 K T = 313.15 K
Sulfadiazine
0.00b 0.0000 0.0380 ± 0.0001 0.0481 ± 0.0002 0.0649 ± 0.0004 0.087 ± 0.002 0.114 ± 0.001
0.10 0.0322 0.085 ± 0.001 0.115 ± 0.001 0.149 ± 0.001 0.198 ± 0.005 0.246 ± 0.006
0.20 0.0697 0.148 ± 0.002 0.200 ± 0.002 0.265 ± 0.001 0.324 ± 0.004 0.424 ± 0.001
0.30 0.1139 0.356 ± 0.001 0.459 ± 0.004 0.598 ± 0.007 0.788 ± 0.013 0.959 ± 0.022
0.40 0.1666 0.482 ± 0.005 0.609 ± 0.002 0.797 ± 0.001 1.016 ± 0.008 1.304 ± 0.011
0.50 0.2306 0.660 ± 0.008 0.808 ± 0.013 1.093 ± 0.014 1.364 ± 0.016 1.739 ± 0.036
0.60 0.3102 0.847 ± 0.002 1.050 ± 0.003 1.364 ± 0.018 1.782 ± 0.037 2.168 ± 0.005
0.70 0.4116 0.958 ± 0.003 1.192 ± 0.006 1.563 ± 0.021 1.871 ± 0.007 2.375 ± 0.028
0.80 0.5453 1.014 ± 0.012 1.233 ± 0.013 1.587 ± 0.014 1.916 ± 0.030 2.408 ± 0.043
0.90 0.7296 0.799 ± 0.007 0.946 ± 0.003 1.235 ± 0.018 1.508 ± 0.004 1.808 ± 0.023
1.00 1.0000 0.388 ± 0.004 0.471 ± 0.006 0.587 ± 0.010 0.721 ± 0.004 0.913 ± 0.012
Idealb 25.45 ± 0.10 30.14 ± 0.12 35.60 ± 0.14 41.92 ± 0.17 49.25 ± 0.20
Sulfamerazine
0.00c 0.0000 0.134 ± 0.002 0.171 ± 0.003 0.209 ± 0.001 0.258 ± 0.003 0.316 ± 0.004
0.10 0.0322 0.281 ± 0.003 0.352 ± 0.004 0.440 ± 0.008 0.545 ± 0.004 0.656 ± 0.012
0.20 0.0697 0.659 ± 0.002 0.864 ± 0.014 1.08 ± 0.01 1.31 ± 0.03 1.63 ± 0.01
0.30 0.1139 1.25 ± 0.02 1.59 ± 0.01 2.05 ± 0.00 2.53 ± 0.02 3.11 ± 0.02
0.40 0.1666 1.85 ± 0.03 2.32 ± 0.02 3.03 ± 0.04 3.68 ± 0.02 4.53 ± 0.06
0.50 0.2306 2.50 ± 0.04 3.24 ± 0.02 4.00 ± 0.06 4.99 ± 0.09 6.12 ± 0.03
0.60 0.3102 3.27 ± 0.01 3.97 ± 0.04 5.01 ± 0.05 6.04 ± 0.06 7.71 ± 0.06
0.70 0.4116 4.12 ± 0.04 4.87 ± 0.05 5.87 ± 0.05 7.12 ± 0.07 8.79 ± 0.27
0.80 0.5453 4.39 ± 0.03 5.11 ± 0.03 6.26 ± 0.05 7.30 ± 0.08 8.62 ± 0.05
0.90 0.7296 3.51 ± 0.01 4.23 ± 0.05 5.02 ± 0.02 5.94 ± 0.05 7.02 ± 0.10
1.00 1.0000 1.78 ± 0.02 2.15 ± 0.02 2.65 ± 0.01 3.07 ± 0.01 3.76 ± 0.00
Idealc 46.2 ± 1.1 54.5 ± 1.3 64.1 ± 1.5 75.3 ± 1.8 88.1 ± 2.1
Sulfamethazine
0.00c 0.0000 0.222 ± 0.001 0.281 ± 0.001 0.367 ± 0.003 0.433 ± 0.001 0.555 ± 0.004
0.10 0.0322 0.594 ± 0.005 0.737 ± 0.001 0.904 ± 0.009 1.12 ± 0.00 1.39 ± 0.01
0.20 0.0697 1.65 ± 0.02 2.01 ± 0.00 2.45 ± 0.03 3.07 ± 0.07 3.84 ± 0.06
0.30 0.1139 3.04 ± 0.09 3.89 ± 0.03 4.50 ± 0.11 5.50 ± 0.04 7.04 ± 0.07
0.40 0.1666 5.05 ± 0.11 5.88 ± 0.06 7.27 ± 0.15 8.42 ± 0.03 11.01 ± 0.15
0.50 0.2306 7.05 ± 0.12 8.41 ± 0.09 9.91 ± 0.03 12.08 ± 0.12 14.69 ± 0.10
0.60 0.3102 9.06 ± 0.13 10.55 ± 0.11 12.71 ± 0.17 15.33 ± 0.13 18.34 ± 0.11
0.70 0.4116 10.86 ± 0.02 12.45 ± 0.11 15.08 ± 0.11 18.43 ± 0.18 21.74 ± 0.20
0.80 0.5453 11.30 ± 0.07 13.73 ± 0.07 16.24 ± 0.16 19.13 ± 0.28 23.18 ± 0.11
0.90 0.7296 10.23 ± 0.30 11.98 ± 0.39 13.77 ± 0.33 17.07 ± 0.19 19.52 ± 0.27
1.00 1.0000 5.91 ± 0.14 6.48 ± 0.22 6.98 ± 0.05 8.40 ± 0.15 10.13 ± 0.13
840 J Solution Chem (2014) 43:836–852
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solute interactions in comparison with the solvent–solute interactions as described later, in
particular in water-rich mixtures.
On the other hand, Fig. 2 shows the solubility profiles of the sulfonamides in 1-pro-
panol ? water and ethanol ? water [18, 19] mixtures as a function of the polarity of the
co-solvent mixtures, expressed by their solubility parameters (dmix), at T = 298.15 K. For
a binary mixture dsolv can be calculated as d1/1 ? d2(1 - /1) from the solubility
parameter of the neat solvents (d1 = 26.5 MPa1/2 for ethanol and 24.4 MPa1/2 for 1-pro-
panol and d2 = 47.8 MPa1/2 for water [24, 25] ) and the volume fraction /i of each
component in the mixture, which is calculated assuming additive volumes [26].
Considering the entire polarity region, all the solubility curves show a maximum at 0.80
mass fraction of 1-propanol (with dsolv = 28.3 MPa1/2), which is similar to that reported in
ethanol ? water mixtures where the maxima were obtained at 0.80 in mass fraction of
ethanol but with dsolv = 30.0 MPa1/2. According to the literature, solutes reach their
maximum solubility in solvents with the same solubility parameter [21] and, therefore, the
d3 values of SD, SMR and SMT should be 28.3 or 30.0 MPa1/2 depending on the co-
solvent considered in the mixtures, i.e. 1-propanol or ethanol, respectively. Nevertheless,
the solubility parameter of SD, SMR and SMT, estimated according to the group contri-
bution method proposed by Fedors [27], are d3 = 28.9 MPa1/2, 28.1 MPa1/2 and
27.4 MPa1/2, respectively [18, 19], which are close to the constant value of maximum
solubilities obtained in 1-propanol ? water mixtures (dsolv = 28.3 MPa1/2).
3.2 Activity Coefficients
Table 4 shows the activity coefficients, c3, of SD, SMR and SMT, calculated as xid3 /x3 from
the respective solubility values presented in Table 2. In almost all cases these values are
similar but slightly higher than the ones exhibited by the same sulfonamides in etha-
nol ? water mixtures [18, 19]. From the c3 values a rough estimate of solute–solvent
intermolecular interactions can be made by considering the following expression [28]:
ln c3 ¼ ðe11 þ e33 � 2e13ÞV3/
21
RTð1Þ
Here subscript 1 stands for the solvent (in the present case, the solvent mixture:
1-propanol ? water), e11, e33 and e13 represent the solvent–solvent, solute–solute and
solvent–solute interaction energies, respectively; V3 is the molar volume of the supercooled
liquid solute, and finally /1 is the volume fraction of the solvent. As a first approximation,
for compounds with low solubility x3, the term V3/12/RT may be considered constant, thus
c3 depends mainly on e11, e33 and e13 [28]. The e11 and e33 terms are unfavorable for
Table 2 continued
w1a x1
a 104 x3
T = 293.15 K T = 298.15 K T = 303.15 K T = 308.15 K T = 313.15 K
Idealc 88.6 ± 1.6 105.0 ± 1.9 124.1 ± 2.2 146.3 ± 2.6 172.0 ± 3.1
a w1 and x1 are the mass and mole fractions of 1-propanol in the co-solvent mixtures free of sulfadiazine,sulfamerazine or sulfamethazine, respectivelyb Values from Ref. [18]c Values from Ref. [19]
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Table 3 Experimental solubility of sulfadiazine, sulfamerazine and sulfamethazine in 1-propanol ? watermixtures expressed in molarity (mol�dm-3) at several temperatures (±0.05 K)
w1a x1
a mol�dm-3 (9 103)
T = 293.15 K T = 298.15 K T = 303.15 K T = 308.15 K T = 313.15 K
Sulfadiazine
0.00b 0.0000 0.211 ± 0.001 0.266 ± 0.001 0.359 ± 0.002 0.479 ± 0.010 0.628 ± 0.008
0.10 0.0322 0.434 ± 0.004 0.581 ± 0.005 0.753 ± 0.003 1.002 ± 0.027 1.239 ± 0.030
0.20 0.0697 0.694 ± 0.012 0.934 ± 0.011 1.238 ± 0.006 1.507 ± 0.020 1.966 ± 0.006
0.30 0.1139 1.485 ± 0.004 1.910 ± 0.015 2.482 ± 0.028 3.263 ± 0.055 3.963 ± 0.089
0.40 0.1666 1.799 ± 0.021 2.273 ± 0.009 2.970 ± 0.003 3.759 ± 0.028 4.791 ± 0.040
0.50 0.2306 2.169 ± 0.025 2.651 ± 0.042 3.579 ± 0.044 4.440 ± 0.051 5.61 ± 0.12
0.60 0.3102 2.425 ± 0.004 2.988 ± 0.010 3.868 ± 0.053 5.03 ± 0.11 6.085 ± 0.013
0.70 0.4116 2.363 ± 0.007 2.916 ± 0.015 3.806 ± 0.051 4.529 ± 0.018 5.722 ± 0.067
0.80 0.5453 2.102 ± 0.025 2.541 ± 0.028 3.259 ± 0.030 3.915 ± 0.061 4.898 ± 0.087
0.90 0.7296 1.356 ± 0.011 1.597 ± 0.006 2.075 ± 0.030 2.521 ± 0.007 3.005 ± 0.038
1.00 1.0000 0.519 ± 0.005 0.628 ± 0.008 0.778 ± 0.013 0.951 ± 0.005 1.198 ± 0.015
Sulfamerazine
0.00c 0.0000 0.74 ± 0.01 0.95 ± 0.01 1.15 ± 0.01 1.42 ± 0.02 1.74 ± 0.02
0.10 0.0322 1.43 ± 0.02 1.79 ± 0.02 2.23 ± 0.04 2.75 ± 0.02 3.31 ± 0.06
0.20 0.0697 3.05 ± 0.01 3.99 ± 0.06 4.95 ± 0.06 6.01 ± 0.12 7.46 ± 0.06
0.30 0.1139 5.19 ± 0.07 6.61 ± 0.03 8.48 ± 0.02 10.48 ± 0.09 12.85 ± 0.07
0.40 0.1666 6.85 ± 0.10 8.58 ± 0.07 11.18 ± 0.16 13.51 ± 0.06 16.57 ± 0.21
0.50 0.2306 8.19 ± 0.13 10.57 ± 0.07 13.01 ± 0.18 16.14 ± 0.30 19.72 ± 0.08
0.60 0.3102 9.34 ± 0.03 11.29 ± 0.10 14.19 ± 0.13 17.00 ± 0.18 21.59 ± 0.16
0.70 0.4116 10.11 ± 0.10 11.91 ± 0.13 14.29 ± 0.13 17.23 ± 0.16 21.17 ± 0.64
0.80 0.5453 9.09 ± 0.06 10.51 ± 0.06 12.79 ± 0.10 14.86 ± 0.17 17.44 ± 0.10
0.90 0.7296 5.95 ± 0.02 7.13 ± 0.08 8.43 ± 0.04 9.93 ± 0.08 11.66 ± 0.16
1.00 1.0000 2.38 ± 0.03 2.87 ± 0.03 3.51 ± 0.01 4.05 ± 0.02 4.94 ± 0.00
Sulfamethazine
0.00c 0.0000 1.23 ± 0.01 1.56 ± 0.01 2.03 ± 0.02 2.39 ± 0.01 3.06 ± 0.02
0.10 0.0322 3.01 ± 0.02 3.74 ± 0.00 4.57 ± 0.05 5.64 ± 0.01 7.01 ± 0.05
0.20 0.0697 7.63 ± 0.09 9.25 ± 0.02 11.25 ± 0.14 14.09 ± 0.33 17.53 ± 0.27
0.30 0.1139 12.68 ± 0.37 16.15 ± 0.13 18.66 ± 0.43 22.70 ± 0.18 28.87 ± 0.30
0.40 0.1666 18.79 ± 0.39 21.80 ± 0.21 26.81 ± 0.53 30.94 ± 0.10 40.24 ± 0.54
0.50 0.2306 23.16 ± 0.38 27.46 ± 0.29 32.17 ± 0.09 39.04 ± 0.41 47.20 ± 0.33
0.60 0.3102 25.87 ± 0.38 29.97 ± 0.31 35.93 ± 0.47 43.07 ± 0.35 51.22 ± 0.29
0.70 0.4116 26.58 ± 0.04 30.29 ± 0.26 36.48 ± 0.26 44.31 ± 0.42 52.01 ± 0.46
0.80 0.5453 23.29 ± 0.14 28.14 ± 0.13 33.04 ± 0.31 38.64 ± 0.56 46.56 ± 0.22
0.90 0.7296 17.29 ± 0.51 20.13 ± 0.65 23.03 ± 0.55 28.36 ± 0.31 32.27 ± 0.45
1.00 1.0000 7.89 ± 0.18 8.62 ± 0.29 9.23 ± 0.07 11.06 ± 0.20 13.25 ± 0.17
a w1 and x1 are the mass and mole fractions of 1-propanol in the co-solvent mixtures free of sulfadiazine,sulfamerazine or sulfamethazine, respectivelyb Values from Ref. [18]c Values from Ref. [19]
842 J Solution Chem (2014) 43:836–852
123
solubility, whereas the e13 term favors the solution process. The contribution of the e33
term can be considered as constant in all the mixtures.
As was described previously for these sulfonamides in ethanol ? water mixtures [18,
19], in a qualitative approach the following analysis can be made based on the energetic
quantities and magnitudes described in Eq. 1. The term e11 is highest in neat water (Hil-
debrand solubility parameter d2 = 47.8 MPa1/2) and is smaller in 1-propanol
(d1 = 24.4 MPa1/2) [23, 24]. Water and water-rich mixtures having larger c3 values (higher
than 260, 150 and 140, for SD, SMR and SMT at 298.15 K for mixtures with w1 = 0.10)
would imply high e11 and low e13 values. On the other hand, in 1-propanol-rich mixtures
(exhibiting c3 values near to 7, 10 and 24 in mixtures with w1 = 0.80 for SD, SMR and
SMT, respectively), the e11 values are relatively low and the e13 values should be relatively
high. Accordingly, the solvation of SD, SMR and SMT should be higher in 1-propanol-rich
mixtures. In a similar way to that described for these compounds in ethanol ? water
mixtures [18, 19], in water and water-rich mixtures the c3 values are highly dependent on
temperature. In all cases the activity coefficients diminish as the temperature rises, giving
more ideal solution processes.
3.3 Apparent Thermodynamic Functions of Solution
The apparent standard enthalpy change of solution is obtained from the Eq. 2 by using the
mean harmonic temperature (Thm) [calculated as: Thm ¼ n=Pn
i¼1
ð1=TÞ], where n is the
number of temperatures studied [29]. Thus, in this case (from 293.15 to 313.15 K) the Thm
value obtained is 303.0 K. In all cases linear weighted regressions were used obtaining
determination coefficients (r2) greater than 0.995 [30].
olnx3
o 1=T � 1=303Kð Þ
� �
p
¼ �DsolnH�
Rð2Þ
The apparent standard Gibbs energy change for the solution process (DsolnG�), con-
sidering the approach proposed by Krug et al. [29], is calculated at 303.0 K by means of:
Fig. 2 Experimental solubility in mole fraction (x3) of sulfadiazine (circles), sulfamerazine (triangles), andsulfamethazine (squares), against the solubility parameter of the 1-propanol ? water (open symbols) andethanol ? water (filled symbols) co-solvent mixtures at 298.15 K
J Solution Chem (2014) 43:836–852 843
123
Table 4 Activity coefficients (c3) of sulfadiazine, sulfamerazine and sulfamethazine in 1-propanol ? watermixtures at several temperatures (±0.05 K)
w1a x1
a T = 293.15 K T = 298.15 K T = 303.15 K T = 308.15 K T = 313.15 K
Sulfadiazine
0.00b 0.0000 670 ± 6 627 ± 6 548 ± 7 483 ± 6 432 ± 7
0.10 0.0322 298 ± 4 263 ± 3 240 ± 2 212 ± 6 200 ± 3
0.20 0.0697 172 ± 3 151 ± 2 134 ± 2 129 ± 2 116 ± 2
0.30 0.1139 71.5 ± 0.7 65.6 ± 0.8 59.5 ± 0.8 53.2 ± 1.0 51.3 ± 0.8
0.40 0.1666 52.8 ± 0.8 49.5 ± 0.5 44.6 ± 0.4 41.3 ± 0.5 37.8 ± 0.6
0.50 0.2306 38.6 ± 0.6 37.3 ± 0.7 32.6 ± 0.5 30.7 ± 0.4 28.3 ± 0.4
0.60 0.3102 30.1 ± 0.3 28.7 ± 0.3 26.1 ± 0.4 23.5 ± 0.5 22.7 ± 0.4
0.70 0.4116 26.6 ± 0.2 25.3 ± 0.3 22.8 ± 0.4 22.4 ± 0.2 20.7 ± 0.3
0.80 0.5453 25.1 ± 0.4 24.4 ± 0.3 22.4 ± 0.3 21.9 ± 0.4 20.5 ± 0.3
0.90 0.7296 31.9 ± 0.4 31.9 ± 0.3 28.8 ± 0.5 27.8 ± 0.3 27.2 ± 0.4
1.00 1.0000 65.7 ± 0.9 64.0 ± 1.0 60.6 ± 1.2 58.1 ± 0.6 53.9 ± 0.8
Sulfamerazine
0.00c 0.0000 345 ± 10 319 ± 9 308 ± 8 292 ± 8 279 ± 7
0.10 0.0322 164 ± 4 154 ± 4 145 ± 4 138 ± 3 134 ± 4
0.20 0.0697 69.8 ± 1.7 62.8 ± 1.8 59.4 ± 1.6 57.3 ± 1.8 53.8 ± 1.3
0.30 0.1139 36.9 ± 1.0 34.1 ± 0.8 31.2 ± 0.7 29.6 ± 0.8 28.2 ± 0.7
0.40 0.1666 24.9 ± 0.7 23.4 ± 0.6 21.1 ± 0.6 20.4 ± 0.5 19.4 ± 0.5
0.50 0.2306 18.4 ± 0.5 16.8 ± 0.4 16.0 ± 0.4 15.0 ± 0.5 14.3 ± 0.3
0.60 0.3102 14.1 ± 0.3 13.7 ± 0.3 12.8 ± 0.3 12.4 ± 0.3 11.4 ± 0.3
0.70 0.4116 11.2 ± 0.3 11.2 ± 0.3 10.9 ± 0.3 10.5 ± 0.3 10.0 ± 0.4
0.80 0.5453 10.5 ± 0.3 10.6 ± 0.3 10.2 ± 0.3 10.3 ± 0.3 10.2 ± 0.2
0.90 0.7296 13.1 ± 0.3 12.9 ± 0.3 12.7 ± 0.3 12.6 ± 0.3 12.5 ± 0.3
1.00 1.0000 25.9 ± 0.7 25.3 ± 0.6 24.1 ± 0.6 24.4 ± 0.6 23.4 ± 0.6
Sulfamethazine
0.00c 0.0000 399 ± 8 373 ± 7 338 ± 7 338 ± 6 310 ± 6
0.10 0.0322 150 ± 3 143 ± 3 138 ± 3 131 ± 2 124 ± 2
0.20 0.0697 53.8 ± 1.2 52.5 ± 1.0 50.8 ± 1.1 47.7 ± 1.4 44.9 ± 1.1
0.30 0.1139 29.2 ± 1.0 27.1 ± 0.5 27.6 ± 0.8 26.6 ± 0.5 24.5 ± 0.5
0.40 0.1666 17.6 ± 0.5 17.9 ± 0.4 17.1 ± 0.5 17.4 ± 0.3 15.7 ± 0.4
0.50 0.2306 12.6 ± 0.3 12.5 ± 0.3 12.6 ± 0.2 12.1 ± 0.3 11.7 ± 0.2
0.60 0.3102 9.8 ± 0.2 10.0 ± 0.2 9.8 ± 0.2 9.6 ± 0.2 9.4 ± 0.2
0.70 0.4116 8.2 ± 0.1 8.5 ± 0.2 8.2 ± 0.2 8.0 ± 0.2 7.9 ± 0.2
0.80 0.5453 7.9 ± 0.1 7.7 ± 0.1 7.7 ± 0.2 7.7 ± 0.2 7.4 ± 0.1
0.90 0.7296 8.7 ± 0.3 8.8 ± 0.3 9.0 ± 0.3 8.6 ± 0.2 8.8 ± 0.2
1.00 1.0000 15.0 ± 0.4 16.2 ± 0.6 17.8 ± 0.3 17.5 ± 0.4 17.0 ± 0.4
a w1 and x1 are the mass and mole fractions of 1-propanol in the co-solvent mixtures free of sulfadiazine,sulfamerazine or sulfamethazine, respectivelyb Values from Ref. [18]c Values from Ref. [19]
844 J Solution Chem (2014) 43:836–852
123
DsolnG� ¼ �R� 303K� intercept ð3Þ
in which the intercept used is the one obtained in the analysis by treatment of ln x3 as a
function of 1/T - 1/Thm. Finally, the standard apparent entropic change for the solution
process (DsolnS�) is obtained from the respective DsolnH� and DsolnG� values at 303.0 K by
using:
DsolnS0 ¼ DsolnH� � DsolnG�ð Þ303K
ð4Þ
Table 5 presents the standard molar thermodynamic functions for dissolution of SD,
SMR and SMT in all of the 1-propanol ? water co-solvent mixtures, including those for
the neat solvents and the ideal solution processes. The propagation of uncertainties in the
thermodynamic quantities calculations was made according to the literature [30, 31].
The standard Gibbs energy of solution is positive in every case as also are the
enthalpy and entropy of solution. Therefore, the dissolution process is always endo-
thermic and entropy driven, according to the increase in entropy after the drug saturates
the solvents. Nevertheless, the thermodynamic function driving the drug’s solubility
increase or decrease according to the co-solvent mixture composition as is shown and
explained later.
In a different way to the behavior reported for SD and SMT in ethanol ? water mix-
tures [18, 19], in 1-propanol ? water mixtures the DsolnH� values of both sulfonamides
diminish from water to 1-propanol-rich mixtures, whereas the DsolnS� values follow a more
erratic behavior. Otherwise, the behavior exhibited by SMR in 1-propanol ? water mix-
tures is similar to the one reported in ethanol ? water mixtures [19]. For this sulfonamide
the DsolnH� and DsolnS� values increase from neat water to w1 = 0.40 or 0.50, but later they
diminish until reaching 1-propanol-rich mixtures.
The relative contributions by enthalpy (fH) and entropy (fTS) toward the solution pro-
cess are given by Eqs. 5 and 6 [32]:
fH ¼DsolnH�j j
DsolnH�j j þ TDsolnS�j j ð5Þ
fTS ¼TDsolnS�j j
DsolnH�j j þ TDsolnS�j j ð6Þ
In all cases the main contributor to the (positive) standard molar Gibbs energy of
solution of both drugs is the positive enthalpy change (fH [ 0.70 for all the sulfonamides),
indicating energetic predominance for the dissolution processes. On the other hand, in a
similar way to that reported in ethanol ? water mixtures [18, 19], the fH values for the
ideal solution processes are lower than the respective values obtained in all of the
experimental solution processes studied. Therefore, the entropy contributions (fTS) are
greater for the ideal dissolution processes, measuring in some way the extent of ran-
domness present in the entropy for the real dissolution processes.
The values of DsolnH� vary nonlinearly with the concentration of 1-propanol in the
aqueous mixtures (Table 5). The addition of 1-propanol to water tends to reduce the
DsolnH� values of SD and SMT solutions at almost all compositions, but it tends to increase
the DsolnH� values of SMR solutions to reach a maximum in the solvent mixture of 0.30 in
mass fraction of 1-propanol. As has been described in the literature, the co-solvent action
in ethanol ? water mixtures may be related to the breaking of the ordered structure of
water (hydrogen bonds) around the non-polar moieties of the drug that increases both the
J Solution Chem (2014) 43:836–852 845
123
Table 5 Apparent thermodynamic functions relative to solution process of sulfadiazine, sulfamerazine andsulfamethazine in 1-propanol ? water mixtures including ideal processes at 303.0 ± 0.05 K
w1a x1
a DsolnG�/kJ�mol-1
DsolnH�/kJ�mol-1
DsolnS�/J�mol-1�K-1
TDsolnS�/kJ�mol-1
fHb fTS
b
Sulfadiazine
0.00c 0.0000 30.0 ± 0.3 43.4 ± 1.1 44.0 ± 1.2 13.3 ± 0.4 0.765 0.235
0.10 0.0322 28.0 ± 0.4 40.7 ± 0.6 41.9 ± 0.9 12.7 ± 0.3 0.762 0.238
0.20 0.0697 26.6 ± 0.3 39.6 ± 0.6 42.6 ± 0.8 12.9 ± 0.3 0.754 0.246
0.30 0.1139 24.5 ± 0.3 38.5 ± 0.6 46.2 ± 0.9 14.0 ± 0.3 0.733 0.267
0.40 0.1666 23.8 ± 0.2 38.2 ± 0.4 47.5 ± 0.5 14.4 ± 0.2 0.726 0.274
0.50 0.2306 23.0 ± 0.3 37.6 ± 0.7 47.9 ± 1.1 14.5 ± 0.3 0.721 0.279
0.60 0.3102 22.4 ± 0.2 36.8 ± 0.6 47.3 ± 0.8 14.4 ± 0.3 0.719 0.281
0.70 0.4116 22.2 ± 0.2 34.6 ± 0.5 41.0 ± 0.7 12.4 ± 0.2 0.736 0.264
0.80 0.5453 22.1 ± 0.3 33.1 ± 0.5 36.4 ± 0.7 11.0 ± 0.2 0.750 0.250
0.90 0.7296 22.7 ± 0.2 32.1 ± 0.6 30.8 ± 0.7 9.3 ± 0.2 0.774 0.226
1.00 1.0000 24.5 ± 0.3 32.6 ± 0.5 26.8 ± 0.5 8.1 ± 0.2 0.801 0.199
Idealc 14.2 ± 0.06 25.2 ± 0.1 36.3 ± 0.3 11.0 ± 0.1 0.696 0.304
Sulfamerazine
0.00d 0.0000 27.2 ± 0.4 32.5 ± 0.4 17.5 ± 0.3 5.3 ± 0.1 0.859 0.141
0.10 0.0322 25.3 ± 0.3 32.6 ± 0.4 24.1 ± 0.4 7.3 ± 0.1 0.817 0.183
0.20 0.0697 23.1 ± 0.3 34.0 ± 0.5 36.0 ± 0.7 10.9 ± 0.2 0.757 0.243
0.30 0.1139 21.5 ± 0.2 35.0 ± 0.3 44.9 ± 0.5 13.6 ± 0.2 0.721 0.279
0.40 0.1666 20.5 ± 0.2 34.5 ± 0.5 46.2 ± 0.8 14.0 ± 0.3 0.711 0.289
0.50 0.2306 19.7 ± 0.2 34.0 ± 0.4 47.0 ± 0.8 14.2 ± 0.2 0.705 0.295
0.60 0.3102 19.2 ± 0.2 32.6 ± 0.5 44.4 ± 0.8 13.5 ± 0.2 0.708 0.292
0.70 0.4116 18.7 ± 0.3 28.9 ± 0.6 33.6 ± 0.9 10.2 ± 0.3 0.739 0.261
0.80 0.5453 18.6 ± 0.1 26.0 ± 0.4 24.4 ± 0.4 7.4 ± 0.1 0.779 0.221
0.90 0.7296 19.2 ± 0.2 26.4 ± 0.2 24.0 ± 0.3 7.3 ± 0.1 0.784 0.216
1.00 1.0000 20.8 ± 0.1 28.3 ± 0.4 24.7 ± 0.4 7.5 ± 0.1 0.791 0.209
Ideald 12.7 ± 0.3 24.6 ± 0.6 39.1 ± 1.3 11.9 ± 0.4 0.675 0.325
Sulfamethazine
0.00d 0.0000 25.8 ± 0.2 34.6 ± 0.6 28.9 ± 0.5 8.8 ± 0.2 0.798 0.202
0.10 0.0322 23.5 ± 0.1 32.4 ± 0.3 29.6 ± 0.3 9.0 ± 0.1 0.783 0.217
0.20 0.0697 20.9 ± 0.3 32.3 ± 0.6 37.5 ± 0.9 11.4 ± 0.3 0.739 0.261
0.30 0.1139 19.4 ± 0.3 30.9 ± 0.9 38.2 ± 1.2 11.6 ± 0.4 0.728 0.272
0.40 0.1666 18.2 ± 0.3 29.2 ± 1.0 36.4 ± 1.3 11.0 ± 0.4 0.726 0.274
0.50 0.2306 17.4 ± 0.2 27.9 ± 0.5 34.7 ± 0.7 10.5 ± 0.2 0.726 0.274
0.60 0.3102 16.8 ± 0.2 27.2 ± 0.5 34.5 ± 0.7 10.5 ± 0.2 0.723 0.277
0.70 0.4116 16.3 ± 0.1 27.1 ± 0.6 35.6 ± 0.8 10.8 ± 0.2 0.716 0.284
0.80 0.5453 16.2 ± 0.1 27.0 ± 0.3 35.7 ± 0.5 10.8 ± 0.2 0.714 0.286
0.90 0.7296 16.5 ± 0.4 25.1 ± 0.9 28.4 ± 1.2 8.6 ± 0.4 0.745 0.255
1.00 1.0000 18.2 ± 0.3 20.3 ± 1.3 7.2 ± 0.5 2.2 ± 0.2 0.903 0.097
846 J Solution Chem (2014) 43:836–852
123
enthalpy of and the entropy of the system [33]. This could be the case for SMR but not for
SD and SMT in 1-propanol ? water mixtures. Above the solvent mixture with 0.40 in
mass fraction of 1-propanol, the apparent enthalpy lowering is the driving force that
enhances the SMR solubility in these media. This behavior is similar to the ones described
previously for this drug in ethanol ? water mixtures [19].
These results indicate that for aqueous-1-propanol mixtures of SMR, two different
driving mechanisms, i.e. entropy or enthalpy, are dominant depending on the co-solvent
composition. At lower 1-propanol ratios (from the water up to the mixtures 0.40 or 0.50 in
mass fraction of 1-propanol the SMR solubility increase is entropy driven, whereas at
greater 1-propanol concentrations (0.50 B w1 B 0.80 or 0.90) the drug solubility
enhancement is enthalpy driven. On the other hand, in the cases of SD and SMT, the
solution process in water rich mixtures is driven by the enthalpy term but the reasons for
this result are unclear.
3.4 Apparent Thermodynamic Functions of Mixing
The dissolution process may be represented by the following hypothetic stages [34],
SoluteðsolidÞ ! SoluteðliquidÞat Tfus ! SoluteðliquidÞat Thm ! SoluteðsolutionÞ
where the dissolution stages are fusion of the solute, cooling of the liquid solute to the
harmonic mean temperature Thm (303.0 K), and the subsequent mixing of the hypothetical
supercooled liquid solute with the solvent at this temperature. This also allows the cal-
culation of the partial thermodynamic contributions to the overall dissolution process by
means of Eqs. 7 and 8, respectively.
DsolnH� ¼ DfusH303 þ DmixH� ð7Þ
DsolnS� ¼ DfusS303 þ DmixS� ð8Þ
where DfusH303 and DfusS
303 represent the thermodynamic functions of fusion of SD, SMR
or SMT and its cooling to the harmonic mean temperature. However, in this research the
DsolnH�-id and DsolnS�-id values for the ideal solution processes were used instead of
DfusH303 and DfusS
303 for reasons described previously in the literature [4]. The same
procedure was used with these sulfonamides in ethanol ? water co-solvent mixtures [18,
19]. Figure 3 summarizes the thermodynamic quantities of mixing of supercooled liquid
SD, SMR and SMT for all of the 1-propanol ? water co-solvent mixtures. The Gibbs
Table 5 continued
w1a x1
a DsolnG�/kJ�mol-1
DsolnH�/kJ�mol-1
DsolnS�/J�mol-1�K-1
TDsolnS�/kJ�mol-1
fHb fTS
b
Ideald 11.1 ± 0.2 25.3 ± 0.5 47.1 ± 1.2 14.3 ± 0.4 0.640 0.360
a w1 and x1 are the mass and mole fractions of 1-propanol in the co-solvent mixtures free of sulfadiazine,sulfamerazine or sulfamethazine, respectivelyb fH and fTS are the relative contributions by enthalpy and entropy towards the Gibbs energy of solution.These values were calculated by means of Eqs. 5 and 6, respectivelyc Values from Ref. [18]d Values from Ref. [19]
J Solution Chem (2014) 43:836–852 847
123
energy of mixing is positive in all cases, which is almost similar to that observed for these
sulfonamides in ethanol ? water mixtures [18, 19] and for sulfanilamide, sulfamethizole
and sulfapyridine in other aqueous alcoholic mixtures [14–17].
The ideal dissolution contributions (related to the solute fusion process) to the enthalpy
and entropy of dissolution of SD, SMR and SMT, DsolnH�-id and DsolnS�-id, are positive
(Table 5). In different way, the contributions of the mixing process toward the overall
dissolution processes are variable, i.e. DmixH� is positive for all drugs in almost all
compositions, except for SMR in 1-propanol rich mixtures, whereas, the entropy of mixing
(DmixS�) is positive for SD in compositions 0.00 B w1 B 0.80 and positive for SMR in
compositions 0.30 B w1 B 0.60 but negative in the other compositions. Finally, DmixS� is
negative for SMT in all the compositions.
According to Fig. 3, the molar DmixG� values diminish as the 1-propanol proportion
increases in the mixtures up to mixture with w1 = 0.80, and they increase in the mixture
with w1 = 0.90 and in neat 1-propanol. The behavior of DmixG� is similar to the one
exhibited by these sulfonamides in ethanol ? water co-solvent mixtures [18, 19]. In a
different way, for SD and SMT the general trend is a nonlinear decrease of DmixH� values
from water up to the mixtures with w1 = 0.90 or neat 1-propanol, respectively. The
enthalpy is responsible for increasing the solubility of SD and SMT in these solvent
mixtures; however, for SD the enthalpy increase from 0.90 mass fraction of 1-propanol up
to neat 1-propanol should corresponds to a decrease of solubility of SD. On the other hand,
the DmixS� values for SD decrease nonlinearly from water up to the mixture with
w1 = 0.10, which opposes the increase in solubility observed, and from this composition
up to 0.50 in mass fraction of 1-propanol the DmixS� increase favors the increase in
solubility of this sulfonamide (SD). With respect to SMR, this sulfonamide exhibits a
similar behavior to SD in the composition range from 0.10 mass fraction of 1-propanol up
to neat 1-propanol. Finally, SMT presents an interesting behavior, i.e. the entropy increases
in water-rich mixtures while the entropy decreases in the 1-propanol-rich mixtures.
However, in intermediate solvent mixtures (0.20 B w1 B 0.80) it remains practically
constant. In a general, the values of the mixing entropy tend to decrease with the increase
of number of methyl groups in the heterocyclic ring of the sulfonamide (SD = 0,
SMR = 1 and SMT = 2, Fig 1), these values being greater for SD and lower for SMT.
Finally, the general behaviors for enthalpy and entropy of mixing of SMR are similar to
those exhibited by the same drugs in ethanol ? water mixtures [18, 19].
As has been previously discussed [14–19], the net variation in DmixH� values (Fig. 3)
results from the contributions of several kinds of interactions. Thus, the enthalpy of cavity
formation (required for the solute accommodation) is endothermic because energy must be
supplied against the cohesive forces of the solvent. This process decreases the drug sol-
ubility, which is in agreement with the discussion of e11 and the solubility parameters of
water and 1-propanol made previously. On the other hand, the enthalpy of solvent–solute
interaction (corresponding to the energy e13) is exothermic and results mainly from van der
Waals and Lewis acid–base interactions. The association of water molecules around the
non-polar groups of the solutes (i.e. hydrophobic hydration) contributes to lowering of the
net DmixH� to small or even negative values in water-rich mixtures. This is not observed in
the case of SD, SMR and SMT in 1-propanol ? water mixtures, just as it was not observed
for the same sulfonamides in ethanol ? water mixtures [18, 19]. Nevertheless, this trend
was reported for sulfapyridine in water [17].
On the other hand, as was previously described [19], by considering the bigger
reduction in entropy obtained with SMR and SMT in neat water in comparison with SD
848 J Solution Chem (2014) 43:836–852
123
(Fig. 3), it is conjecturable that the hydrophobic hydration around the methyl-substituted
heterocyclic ring will be greater in SMR and SMT than in the heterocyclic ring of SD.
It is expected that the energy of cavity formation should be lower as the proportion of
1-propanol increases. This effect is well observed for both SMR and SMT in 1-propanol-
Fig. 3 Apparent thermodynamic quantities of mixing of sulfadiazine (open circle), sulfamerazine (opentriangle), and sulfamethazine (open square), in 1-propanol ? water mixtures at 303.0 K as a function of theco-solvent mixture composition
J Solution Chem (2014) 43:836–852 849
123
rich mixtures (w1 C 0.20 and w1 C 0.30, respectively), where the DmixH� values diminish
as the proportion of co-solvent increases. According to Romero et al. [33], in the initial
portion of the solubility curve the hydrogen bonding of the drug will increase with the co-
solvent concentration in the co-solvent mixtures. However, at large co-solvent proportions
this kind of interaction may become saturated, thus becoming a constant contribution.
Otherwise, nonspecific and cavity effects are not saturated and can vary with the co-solvent
concentration. Nevertheless, these results are observed only for SMR, whereas for SD and
SMT a continuous decrease in mixing enthalpy is observed from neat water up to the
mixture with 0.90 mass fraction of 1-propanol, but then increase slightly in neat 1-pro-
panol. Thus these results for the last two sulfonamides are different with respect to those
exhibited by the same drugs in ethanol ? water mixtures [18, 19]. Nevertheless, the rea-
sons for these behaviors are unclear but they could be associated to the higher hydrophobic
hydration of the propyl moiety of this co-solvent in comparison with ethanol.
3.5 Enthalpy–Entropy Compensation Analysis
There are some reports in the literature demonstrating non-enthalpy–entropy compensation
in the solubility of drugs in some aqueous co-solvent mixtures. These analyses have been
used in order to identify the mechanism of the co-solvent’s action. Thus, weighted fits of
DsolnH� as a function of DsolnG� or DmixH� as a function of DmixG� at the harmonic mean
temperature permit such an analysis [19, 35].
Figure 4 shows that SD, SMR and SMT in the 1-propanol ? water co-solvent system
present nonlinear DmixH� versus DmixG� curves with variable positive slope in the interval
from water up to the mixture with w1 = 0.80 for SD and SMT. Beyond this 1-propanol
proportion, up to w1 = 0.90 for SD and up to neat 1-propanol for SMT, negative slopes are
obtained for both. Otherwise, for SMR a variable but positive slope is observed from water
up to the mixture with w1 = 0.30. In the composition range 0.30 B w1 B 0.80, a variable
negative slope is found, and finally, in 1-propanol-rich mixtures negative slopes are also
observed. Accordingly, the mechanism for the dissolution of SMR in water-rich mixtures is
the entropy driven one, probably implying loosening of the water structure, whereas in the
Fig. 4 DmixH� versus DmixG� enthalpy–entropy compensation plot for dissolution process of sulfadiazine(open circle), sulfamerazine (open triangle), and sulfamethazine (open square), in 1-propanol ? water co-solvent mixtures at 303.0 K
850 J Solution Chem (2014) 43:836–852
123
cases with positive slopes the mechanisms is the enthalpy driven, probably due to better
solvation of the drugs by the 1-propanol molecules.
4 Conclusions
From all topics discussed here it can be concluded that the solution process of SD, SMR
and SMT in 1-propanol ? water mixtures depends strongly on the solvent composition as
was also observed for these drugs in ethanol ? water mixtures [18, 19] and for sulfanil-
amide, sulfamethizole and sulfapyridine in similar aqueous alcoholic solutions [14–17].
Nonlinear enthalpy–entropy compensations were found for these drugs in this co-solvent
system. In this context, the dissolution process in water-rich mixtures was found to be
entropy driven only for SMR, whereas for the other two drugs it is found to be enthalpy
driven in almost all of the mixtures. Ultimately, it can be said that the data presented in this
report amplify the physicochemical information about sulfonamide drugs in binary aque-
ous-co-solvent mixtures.
Acknowledgments We thank the Department of Pharmacy of the National University of Colombia forfacilitating us the equipment and laboratories used.
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