Accepted Manuscript

Title: Solubility and solution thermodynamics ofsulfamerazine and sulfamethazine in some ethanol + watermixtures

Author: Daniel R. Delgado Fleming Martı́nez

PII: S0378-3812(13)00516-5DOI: http://dx.doi.org/doi:10.1016/j.fluid.2013.09.018Reference: FLUID 9766

To appear in: Fluid Phase Equilibria

Received date: 10-7-2013Revised date: 3-8-2013Accepted date: 11-9-2013

Please cite this article as: D.R. Delgado, F. Martı́nez, Solubility and solutionthermodynamics of sulfamerazine and sulfamethazine in some ethanol + water mixtures,<i>Fluid Phase Equilibria</i> (2013), http://dx.doi.org/10.1016/j.fluid.2013.09.018

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Solubility and solution thermodynamics of sulfamerazine and sulfamethazine in some ethanol +

water mixtures

Daniel R. Delgado 1, Fleming Martínez 1*

1 Grupo de Investigaciones Farmacéutico-Fisicoquímicas, Departamento de Farmacia, Universidad

Nacional de Colombia, A.A. 14490, Bogotá D.C., Colombia.

* Corresponding author. Tel.: +571 3165000x14608; fax: +571 3165060. E-mail address:

fmartinezr@unal.edu.co (F. Martínez).

Abstract

The solubility of sulfamerazine (SMR) and sulfamethazine (SMT) in some ethanol + water cosolvent

mixtures was measured at five temperatures from 293.15 to 313.15 K in all the polarity range provided by

the aqueous mixtures. The mole fraction solubility of both drugs was maximal in the mixture 0.80 in mass

fraction of ethanol (δ = 30.0 MPa1/2) and minimum in pure water (δ = 47.8 MPa1/2) at all the temperatures

studied. The thermodynamic functions Gibbs energy, enthalpy, and entropy of solution were obtained

from these solubility data by using the van’t Hoff and Gibbs equations. Thermodynamic quantities of

mixing were also calculated by using some calorimetric values related to the drugs fusion process

reported in the literature. Non-linear enthalpy–entropy relationships were observed for both drugs in the

plot of enthalpy vs. Gibbs energy of mixing. The plot of mixH° vs. mixG° shows three different trends

according to the slopes obtained when the mixtures composition changes. Accordingly, the driving

mechanism for SMR solution process in water-rich is the entropy; whereas, from 0.20 mass fraction of

ethanol to neat ethanol the process is enthalpy-driven. For SMT the process is driving by entropy in

water-rich and ethanol-rich mixtures but driven by enthalpy in mixtures from 0.30 to 0.80 mass fraction

of ethanol. The behavior of SMT is similar to the ones exhibited by the similar drugs sulfapyridine and

sulfadiazine in the same co-solvent mixtures.

Key words: sulfamerazine, sulfamethazine, ethanol + water mixtures, solubility, solution

thermodynamics, activity coefficient.

1 Introduction

The behavior of drugs in several co-solvent mixtures as a function of temperature is mainly evaluated for

the purposes of substances purification, liquid medicines design, and understanding of the mechanisms

involved in the physical and chemical stability of pharmaceutical dissolutions [1]. Therefore, solubility is

a very important physicochemical property for pharmaceutical product design because it affects the drug

efficacy, influencing several biopharmaceutical and pharmacokinetic properties [2]. On the other hand,

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temperature-dependence of the solubility allows performing a thermodynamic analysis in order to insight

into the molecular mechanisms involved in the drug dissolution processes [3].

Sulfamerazine (SMR, molar mass 264.31 g mol–1, Fig. 1) and sulfamethazine (SMT, molar mass

278.33 g mol–1, Fig. 1) are two sulfonamide drugs used as effective antimicrobial agents for the

prevention and cure of several kinds of bacterial infections in human and veterinary therapies [4].

***Fig. 1***

Although SMR and SMT have been widely used in therapeutics, the solubility data of these drugs in

co-solvent mixtures is still scarce in the literature [5]. Although some theoretical and semiempirical

models can be used to predict drug solubilities, the availability of experimental data is still fundamental

for the pharmaceutical scientists [6]. Because the solubility of sulfonamides in neat water is too low [7,

8], some co-solvent + water mixtures have been evaluated in order to increase the solubility of several of

these compounds [9, 10]. This has also been made with the purpose to understand the molecular

mechanisms involved in the drug dissolution processes.

According to the literature, ethanol has been studied in particular as possible co-solvent in the design

of several medicines intended for oral and parenteral administration [1, 11]. Moreover, ethanol has also

been used as antimicrobial and/or flavoring agent in several liquids formulations [12]. This co-solvent is a

hydrogen-donor and hydrogen-acceptor compound due to its hydroxyl group, and thus, it is miscible with

water in all proportions [13].

For all these reasons, the main goal of this work is thus to extend the database on experimental solubility

for 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 conformed by ethanol and water, based on the van’t Hoff

method, including the respective contributions by mixing of this compound toward the solution processes, as

has been made with other sulfonamides in other co-solvent systems [14-16]. This thermodynamic study is

very similar to the ones reported previously about the solubility of sulfapyridine [17] and sulfadiazine [18] in

the same ethanol + water mixtures. Sulfapyridine is similar to sulfadiazine but differs from this because it has

only one nitrogen atom in the heterocyclic ring (Fig. 1).

2 Experimental

2.1 Materials

The solutes sulfamerazine (SMR, 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 Co., the absolute ethanol A.R. from Merck

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(EtOH, component 1, with purity greater than 0.995 in mass fraction) and the distilled water (component

2) with conductivity < 2 μS cm–1, were used in agreement with the quality requirements of the American

Pharmacopeia, USP [19]. Molecular sieve (Merck, number 3, pore size 0.3 nm) and Durapore® filters

(0.45 µm, 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.

***Table 1***

2.2 Solvent mixtures preparation

All ethanol (1) + water (2) 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 all

compositions range, the mass fractions of ethanol, w1, of the nine binary mixtures prepared varied by 0.10

from 0.10 to 0.90.

2.3 Solubility determinations

The procedures used in this research were similar to those employed previously in the study of

sulfapyridine and sulfadiazine in the same ethanol + water mixtures [17, 18]. Briefly, an excess of SMR

(3) or SMT (3) 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® E 60

H Elmasonic) during 15 min and later they were placed in thermostatic mechanical shakers (Julabo

SW23) kept at 303.15, 308.15, or 313.15 ( 0.05) K or placed in re-circulating thermostatic baths (Neslab

RTE 10 Digital One Thermo Electron Company) kept at 293.15 or 298.15 ( 0.05) K for at least four

days to reach the equilibrium. This equilibrium time was established by measuring the drug concentration

in neat water at 293.15 K till they became constant. After this time the supernatant solutions were filtered

at isothermal conditions (Millipore Corp. Swinnex®-13) 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 both drugs (UV/VIS BioMate 3 Thermo Electron Company spectrophotometer)

and interpolation from previously constructed UV spectrophotometric calibration curves. All the

solubility experiments were run at least in triplicates. In order to transform mole fractions to molar

concentrations (mol dm–3), the density of the saturated solutions was determined by using a digital density

meter (DMA 45 Anton Paar) connected to the same re-circulating thermostatic baths according to

procedures described in the literature [20].

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3 Results and discussion

In order to propose the possible intermolecular interactions present in the saturated solutions of SMR

(3) or SMT (3), it is important to keep in mind that these drugs, in similar way to sulfapyridine and

sulfadiazine, act in solution mainly as a Lewis bases (due to their –NH2, –SO2–, and =N– groups) and as a

Lewis acids (due to their –NH2 and >N–H groups) in order to establish hydrogen bonds with the –OH

groups in the solvents [17, 18, 21].

3.1 Experimental and ideal solubility

Tables 2 and 3 list the experimental solubility (expressed in mole fraction and molarity, respectively)

of SMR (3) and SMT (3) in EtOH (1) + water (2) mixtures at the temperature range studied, 293.15 to

313.15 K. This temperature range covers different room conditions and also the normal human body

temperature. In almost all cases the variation coefficients of the solubility for both drugs were smaller

than 2.0 %.

For SMR (3) the mole fraction solubility values in neat water (2) at temperatures from 298.15 to

313.15 K are almost 20% lower than the ones reported by Martínez and Gómez (i.e. 1.450 x 10–5 at

298.15 K, 1.724 x 10–5 at 303.15 K, 2.242 x 10–5 at 308.15 K, and 2.816 x 10–5 at 313.15 K, respectively)

[22]. The solubility values for SMT are almost similar with the ones reported in the same reference,

except at 313.15 K (i.e. 2.896 x 10–5 at 298.15 K, 3.613 x 10–5 at 303.15 K, 4.244 x 10–5 at 308.15 K, and

5.175 x 10–5 at 313.15 K, respectively) [22]. Nevertheless, it is important to note that the values reported

by these authors were determined in buffers with ionic strength of 0.15 mol dm–3 instead of neat water

(1).

On the other hand, some differences varying from 25 to 45% are found between data reported by

Zhang et al. and data reported in Table 2 about the mole fraction solubility of SMT (3) in neat ethanol (1)

at all the temperatures studied here (i.e. 9.410 x 10–5 at 293.15 K, 1.216 x 10–4 at 298.15 K, 1.516 x 10–4 at

303.15 K, 1.925 x 10–4 at 308.15 K, and 2.383 x 10–4 at 313.15 K, respectively) [23]. Finally, up to the

best of our knowledge, no solubility values have been published for SMR (3) in ethanol (1) or for both

drugs in ethanol (1) + water (2) mixtures, and therefore, no other comparison is possible.

Solubility increases with temperature in all cases indicating that the dissolution process is

endothermic. The highest mole fraction solubilities of SMR (3) and SMT (3) were obtained in the mixture

0.80 in mass fraction of ethanol at T = 313.15 K, whereas the lowest values were found in pure water (2)

at 293.15 K (Table 2). Nevertheless, if the molarity concentration scale is considered the maximum

solubilities are obtained in the mixture 0.70 in mass fraction of ethanol (Table 3). All these behaviors are

similar to the ones exhibited by sulfapyridine and sulfadiazine in the same co-solvent mixtures [17, 18].

***Tables 2 and 3***

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Table 2 also shows the ideal solubilities in mole fraction of the solutes ( id3x ) calculated by using Eq.

1 with the thermal values of drugs fusion which were taken from the literature, i.e. Tfus = 508.5 K and

ΔfusH = 41.3 ± 1.0 kJ mol–1 for SMR (3) and Tfus = 469.0 K and ΔfusH = 39.2 ± 0.7 kJ mol–1 for SMT (3)

[22].

fus

fusp

fus

fusfusid3 ln

)()(ln

T

T

T

TT

R

C

TRT

TTHx (1)

where id3x is the ideal solubility of the solute as mole fraction, fusH is the molar enthalpy of fusion of the

pure solute (at the melting point), Tfus is the absolute melting point, T is the absolute solution temperature,

R is the gas constant, and Cp is the difference between the molar heat capacity of the crystalline form

and the molar heat capacity of the hypothetical super-cooled liquid form, both at the solution temperature

[24]. Since Cp values are not easily available in the literature it is usual assuming that it may be

approximated to the entropy of fusion, fusS calculated as the quotient fusH/Tfus (i.e. 81.2 ± 1.9 J mol–1 K–

1 for SMR (3) and 83.6 ± 1.5 J mol–1 K–1 for SMT (3), respectively [22]). Table 2 shows that the ideal

solubilities of SMR and SMT are greater than the experimental solubilities obtained at all the

temperatures studied.

On the other hand Figures 2 and 3 shows the solubility profiles, as a function of the polarity of the

mixtures, expressed by their solubility parameters (δmix). For a binary mixture δmix can be calculated from

the solubility parameter of the pure solvents (δ = 26.5 MPa1/2 for ethanol and δ = 47.8 MPa1/2 for water

[25, 26]) and the volume fraction i of each component in the mixture, which is calculated assuming

additive volumes [27, 28]:

n

iii

1mix (2)

***Figs. 2 y 3***

Considering the entire polarity region both solubility curves show a maximum at 0.80 in mass

fraction of ethanol (1) (δmix = 30.0 MPa1/2). In similar way to described previously for sulfapyridine and

sulfadiazine [17, 18], at the left part of the maximum (δmix < 30.0 MPa1/2) the solubility of SMR (3) and

SMT (3) decreases as the solubility parameter of the mixtures is lowered, i.e. from δmix = 30.0 MPa1/2 to

26.5 MPa1/2 (from 0.80 in mass fraction of ethanol (1) to neat ethanol (1)). According to the literature,

solutes reach their maximum solubility in solvents with the same solubility parameter [21] and therefore,

the δ values of SMR (3) and SMT (3) as well as those for sulfapyridine and sulfadiazine would be 30.0

MPa1/2. Nevertheless, the solubility parameter of SMR (3) and SMT (3), estimated according to the

groups contribution method proposed by Fedors [26, 29], are δ = 28.1 MPa1/2 and 27.4 MPa1/2,

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respectively (Table 4), which are lower than the experimental values obtained in this work at the

solubility maximums (δ = 30.0 MPa1/2). Similar results were also observed with sulfapyridine and

sulfadiazine [17, 18]. Thus, this fact indicates that the actual polarity of all these drugs is higher than the

ones expected from the additive contribution of its groups.

***Table 4***

3.2 Activity coefficients

Table 5 shows the activity coefficients of SMR (3) and SMT (3) 3, calculated as id3x /x3 from the

respective solubility values presented in Table 2. In all cases these values are similar to the ones exhibited

by sulfapyridine [17] but lower than the ones exhibited by sulfadiazine in the same mixtures [18]. From 3

values a rough estimate of solute-solvent intermolecular interactions can be made by considering the

following expression [30]:

RT

Veee

213

1333113 )2(ln

(3)

Here subscript 1 stands for the solvent in the present case, the solvent mixture: ethanol (1) + water (2)),

e11, e33 and e13 represent the solvent-solvent, solute-solute and solvent-solute interaction energies,

respectively; V3 is the molar volume of the super-cooled liquid solute, and finally, φ1 is the volume

fraction of the solvent. As a first approximation, for compounds with low solubility x3, the term V3φ12/RT

may be considered constant; thus, γ3 depends mainly on e11, e33 and e13 [30, 31]. The e11 and e33 terms are

unfavorable for solubility, whereas the e13 term favors the solution process. The contribution of the e33

term could be considered as constant in all the mixtures.

***Table 5***

As was described previously for other sulfonamides [17, 18], in a qualitative approach, the following

analysis could be made based on the energetic quantities and magnitudes described in the Eq. (3): The

term e11 is highest in neat water (2) (Hildebrand solubility parameter = 47.8 MPa1/2) and is smaller in

ethanol (1) ( = 26.5 MPa1/2) [28, 29]. Pure water (2) and water-rich mixtures having larger 3 values

(even higher than 300) would imply high e11 and low e13 values. On the other hand, in ethanol-rich

mixtures (having 3 values between 10 and 17), the e11 values are relatively low and the e13 values would

relatively be high. Accordingly, the solvation of SMR (3) and SMT (3) could be just higher in ethanol-

rich mixtures. In water and water-rich mixtures the 3 values are highly dependent on temperature. In all

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cases the activity coefficients diminish as the temperature raises being more ideal the solution processes.

This behavior is almost similar to those obtained with sulfapyridine and sulfadiazine [17, 18].

3.3 Thermodynamic functions of solution

Apparent standard enthalpy change of solution is obtained from the Eq. (4) by using the mean

harmonic temperature (Thm) [calculated as:

n

i

TnT1

hm )/1(/ ], where n is the number of temperatures

studied [32]. Thus, in this case (from 293.15 K 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

[33].

R

H

T

x

P

soln3

K303/1/1

ln(4)

The apparent standard Gibbs energy change for the solution process (∆solnG°), considering the

approach proposed by Krug et al. [32], is calculated at 303.0 K by means of:

interceptK303soln RG (5)

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 solution process (∆solnS°) is obtained from the

respective ∆solnH° and ∆solnG° values at 303.0 K by using:

K303

solnsoln0soln

GHS (6)

Table 6 presents the standard molar thermodynamic functions for dissolution of SMR (3) and SMT

(3) in all the ethanol (1) + water (2) 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 [33, 34].

***Table 6***

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 endothermic and entropy-driven. In general way,

the ΔsolnH° and ΔsolnS° values increase from pure water (2) to the mixture w1 = 0.20 for SMR (3) or to the

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mixture w1 = 0.30 for SMT (3) and later they diminish with the ethanol proportion to the mixture w1 =

0.90 for SMR (3) and to the pure ethanol (1) for SMT (3).

The relative contributions by enthalpy (H) and entropy (TS) toward the solution process are given

by equations (7) and (8) [35].

STH

HH

solnsoln

soln (7)

STH

STTS

solnsoln

soln (8)

In all cases the main contributor to the (positive) standard molar Gibbs energy of solution of both

drugs is the positive enthalpy (H > 0.70 for SMR (3) and H > 0.65 for SMT (3)), indicating energetic

predominance on the dissolution processes. In similar way to obtained with sulfapyridine and sulfadiazine

in the same co-solvent mixtures [17, 18], the minimum enthalpy-contribution for SMR (3) is obtained in

the mixture with w1 = 0.30. In the case of SMT (3) this value is obtained in the mixture with w1 = 0.60.

On the other hand, H values for the ideal solution processes are lower than the respective values obtained

in all the experimental solution processes studied. Therefore, the entropy contributions (TS) are greater

for the ideal processes indicating in some way the extent of restriction presented in entropy in real

solution processes.

The values of ΔsolnH° vary non-linearly with the ethanol concentration in the aqueous mixtures

(Table 6). The addition of ethanol (1) to water (2) tends to increase the ΔsolnH° values of SMR (3) and

SMT (3) to a maximum in the mixtures 0.20 and 0.30 in mass fraction of ethanol (1), respectively. As has

been described in the literature, the co-solvent action 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

enthalpy of and the entropy of the system [36]. Above the mixtures 0.20 and 0.30 in mass fraction of

ethanol (1), the apparent enthalpy lowering is the driving force that enhances the drug solubility in both

cases. This behavior is similar to the ones described previously for sulfapyridine and sulfadiazine in the

same co-solvent mixtures [17, 18].

These results indicate that for the aqueous mixtures two different driving-mechanisms; thus, entropy

or enthalpy are dominant depending on the co-solvent composition. At lower ethanol (1) ratios (from the

pure water (2) up to the mixtures 0.20 or 0.30 in mass fraction of ethanol (1)) SMR (3) and SMT (3)

solubility increase is entropy-driven whereas at greater ethanol (1) concentrations (0.20 w1 0.90 for

SMR (3) and 0.20 w1 1.00 for SMT (3)) the drugs solubility enhancement is enthalpy-driven. In the

case of SMR (3) from the mixture 0.90 in mass fraction of ethanol (1) to neat ethanol (1) the solution

process is driven by entropy as will be discussed later.

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3.4 Thermodynamic functions of mixing

The dissolution process may be represented by the following hypothetic stages [37],

Solute(Solid) Solute(Liquid) at Tfus Solute(Liquid) at Thm Solute(Solution)

where the dissolution stages are the solute fusion, the cooling the liquid solute to the harmonic mean

temperature Thm (303.0 K), and the subsequent mixing of the hypothetical super-cooled liquid solute with

the solvent at this temperature. This allows also the calculation of the partial thermodynamic

contributions to the overall dissolution process by means of equations (9) and (10), respectively.

ΔsolnH° = ΔfusH303 + ΔmixH° (9)

ΔsolnS° = ΔfusS303 + ΔmixS° (10)

where ΔfusH303 and ΔfusS

303 represent the thermodynamic functions of fusion of SMR (3) or SMT (3) and

its cooling to the harmonic mean temperature. However, in this research the ΔsolnH°-id and ΔsolnS°-id values

for the ideal solution processes were used instead of ΔfusH303 and ΔfusS

303 for reasons described previously

in the literature [3]. The same procedure was used with sulfapyridine and sulfadiazine in the same co-

solvent system [17, 18]. Fig. 4 summarizes the thermodynamic quantities of mixing of super-cooled

liquid SMR (3) and SMT (3) with all the co-solvent mixtures. Gibbs energy of mixing is positive in all

cases, which is almost similar to that observed for sulfapyridine and sulfadiazine in ethanol (1) + water

(2) mixtures [17, 18] and for sulfanilamide, sulfamethizole and sulfapyridine in propylene glycol + water

mixtures [14-16].

***Fig. 4***

The ideal dissolution contributions (related to solute fusion process) to the enthalpy and entropy of

dissolution of SMR (3) and SMT (3), ΔsolnH°-id and ΔsolnS°-id, are positive (Table 6). The contribution of

the mixing process toward the overall dissolution is variable: ΔmixH° is positive for both drugs in all

compositions; whereas, the entropy of mixing (ΔmixS°) is positive for SMR (3) in mixtures with 0.10 w1

0.70 and for SMT (3) in mixtures with 0.20 w1 0.80 and it is negative in the other mixtures.

According to Fig. 4, the molar ΔmixG° values diminish as the ethanol (1) proportion increases in the

mixtures up to mixture with w1 = 0.80 and they increases in the mixture with w1 = 0.90 and in neat

ethanol (1); whereas, the ΔmixH° values increase nonlinearly from pure water (2) up to the mixtures with

w1 = 0.20 or 0.30 for SMR (3) and SMT (3) respectively, where the greatest enthalpic values are also

obtained and later they diminish. The behavior of ΔmixG° is just similar to the one exhibited by

sulfapyridine and sulfadiazine in the same co-solvent mixtures [17, 18]. On the other hand, in similar way

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to the enthalpy of mixing, the ΔmixS° values increase nonlinearly from pure water (2) up to the mixture

with w1 = 0.20 or 0.40 for SMR (3) and SMT (3) respectively, and later they diminish. The general

behaviors for enthalpy and entropy of mixing of SMT (3) are also similar to the ones exhibited by

sulfapyridine and sulfapyridine in the same mixtures [17, 18].

The net variation in ΔmixH° values (Fig. 4) results from the contribution of several kinds of

interactions. Thus, the enthalpy of cavity formation (required for solute accommodation) is endothermic

because energy must be supplied against the cohesive forces of the solvent. This process decreases the

drug solubility, which is in agreement with the discussion of e11 and the solubility parameters of water (2)

and of ethanol (1) 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 structuring of water molecules around the non-polar groups of the solutes (i.e.

hydrophobic hydration) contributes to lowering of the net ΔmixH° to small or even negative values in

water-rich mixtures. This fact is not observed in the case of SMR (3) and SMT (3) as it was not observed

for sulfadiazine [18] but it was reported for sulfapyridine in pure water [17].

On the other hand, by considering the bigger reduction in entropy obtained with SMR (3) and SMT

(3) in neat water in comparison with sulfadiazine (Fig. 4), it is conjecturable that the hydrophobic

hydration around the methyl-substituted heterocyclic ring would be greater in SMR (3) and SMT (3) than

in the heterocyclic ring of sulfadiazine.

The energy of cavity formation should be lower as the proportion of ethanol (1) increases. This effect

is well observed for both SMR (3) and SMT (3) in ethanol-rich mixtures (w1 ≥ 0.20 and w1 ≥ 0.30,

respectively), where ΔmixH° diminish as the proportion of co-solvent increases. According to Romero et

al. [36] in the initial portion of the solubility curve the hydrogen bonding of the drug will increase with

ethanol (1) concentration in the co-solvent mixtures, just as also occurs with sulfapyridine [17] and

sulfadiazine [18], and with the analgesic drugs acetaminophen [38], naproxen [3], ketoprofen [39], and

indomethacin [28] in the same co-solvent system. However, at large co-solvent proportions this

interaction may be saturated, becoming a constant contribution. On the other hand, nonspecific and cavity

effects are not saturated and vary with co-solvent concentration. Although, it is clear that for SMR (3) and

SMT (3) the values of ΔmixH° also diminish as the proportion of ethanol (1) increases in the mixtures, as

was already said.

3.5 Thermodynamic functions of transfer

In order to verify the effect of co-solvent composition on the thermodynamic function driving the

solution process, Table 7 summarizes the thermodynamic functions of transfer of SMR (3) and SMT (3)

from the more polar solvents to the less polar ones. These new functions were calculated as the

differences between the thermodynamic quantities of solution obtained in the less polar mixtures and the

more polar ones. This procedure is the same followed previously with sulfapyridine and sulfadiazine in

the same co-solvent mixtures [17, 18].

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If the addition of ethanol (1) to water (2) is considered [being the solvent mixture less polar as the

ethanol (1) proportion increases], as has been done earlier with sulfapyridine and sulfadiazine [17, 18]

and some other drugs [3, 24, 38, 39], it happens the following for SMR: from neat water (2) to 0.20 in

mass fraction of ethanol (1) (∆A→BG° < 0, ∆A→BH° > 0, and ∆A→BS° > 0) the dissolution process is driven

by the entropy. From 0.20 to 0.80 in mass fraction of ethanol (1) (∆A→BG° < 0, ∆A→BH° < 0, and ∆A→BS°

< 0) the dissolution process is enthalpy driven. From 0.80 to 0.90 in mass fraction of ethanol (1) (∆A→BG°

> 0, ∆A→BH° > 0, and ∆A→BS° < 0) the dissolution process is apparently enthalpy and entropy driven.

Ultimately, from this ethanol (1) proportion to neat ethanol (1) (∆A→BG° > 0, ∆A→BH° > 0, and ∆A→BS° >

0) the dissolution process is enthalpy driven again. All these behaviors are almost similar to the ones

exhibited by sulfapyridine and sulfadiazine in the same co-solvent mixtures [17, 18]. In the case of SMT

happens the following: from neat water (2) to 0.30 in mass fraction of ethanol (1) (∆A→BG° < 0, ∆A→BH° >

0, and ∆A→BS° > 0) the dissolution process is driven by the entropy. From 0.30 to 0.40 in mass fraction of

ethanol (1) (∆A→BG° < 0, ∆A→BH° < 0, and ∆A→BS° > 0) the dissolution process is enthalpy and entropy

driven. From 0.40 to 0.80 in mass fraction of ethanol (1) (∆A→BG° < 0, ∆A→BH° < 0, and ∆A→BS° < 0) the

dissolution process is enthalpy driven. Ultimately, from this ethanol (1) proportion to neat ethanol (1)

(∆A→BG° > 0, ∆A→BH° < 0, and ∆A→BS° < 0) the dissolution process is entropy driven again. These results

could be interpreted as the result of the water-structure losing around the non-polar groups of the drugs

due to the addition of co-solvent in the water-rich mixtures and to more drugs solvation by co-solvent

molecules in the co-solvent-rich mixtures.

***Table 7***

3.6 Enthalpy-entropy compensation analysis

There are some reports in the literature demonstrating non-enthalpy-entropy compensation in the

solubility of drugs in several aqueous co-solvent mixtures. These analyses have been used in order to

identify the mechanism of the co-solvent action. Thus, weighted graphs of ΔsolnH° as a function of ΔsolnG°

or ΔmixH° as a function of ΔmixG° at the harmonic mean temperature permit such an analysis [18, 40].

Fig. 5 shows that SMR (3) and SMT (3) in the ethanol (1) + water (2) co-solvent system present

non-linear ΔmixH° vs. ΔmixG° curves with a variable negative slope in the interval from pure water (2) up

to w1 = 0.20 for SMR (3) and from pure water (2) up to w1 = 0.30 for SMT (3). Beyond this ethanol (1)

proportions up to w1 = 0.80 a variable but positive slope is obtained is obtained for both drugs, and

finally, from this mixture to pure ethanol (1) a negative slope is obtained again for SMT (3) and a positive

but no continuous slope for SMR (3). As can be seen in Fig. 5, the behavior of SMT (3) is similar to the

one exhibited by sulfadiazine. Accordingly, the driving mechanism for the dissolution of SMR (3) and

SMT (3) is the entropy in the case with negative slope probably implying water-structure loosening,

whereas in the case with positive slope the driving mechanism is the enthalpy, probably due to better

solvation of the drug by ethanol (1) molecules, as was already said.

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***Fig. 5***

4 Conclusions

From all topics discussed here it can be concluded that the solution process of SMR (3) and SMT (3) in

ethanol (1) + water (2) mixtures depends strongly on the solvent composition as was also observed for

sulfapyridine and sulfadiazine in the same cosolvent mixtures [17, 18] and for sulfanilamide,

sulfamethizole and sulfapyridine in propylene glycol + water mixtures [14-16]. In similar way to

observed with sulfapyridine and sulfadiazine, the hydrogen bonding because of the sulfonamide and

amine groups of SMR (3) and SMT (3) to the more basic solvent component, i.e. ethanol (1), could to

cause the latter to solvate the drug molecules preferentially [37, 41]. Nevertheless, this is just a part of the

cause for the higher solubility of these sulfonamides in the ethanol-rich mixtures compared with the

water-rich ones, being the other reason the strong self-interaction of the water (2) molecules that hinders

the introduction of these bulky solutes into the solutions. Non-linear enthalpy-entropy compensations

were found for these drugs in this solvent system. In this context, entropy-driving was found for the drugs

dissolution processes in water-rich mixtures and enthalpy-driving in the other mixtures. Nevertheless,

entropy-driving is found for SMT (3) in ethanol-rich mixtures. Ultimately, it can be said that the data

presented in this report amplify the physicochemical information about anti-parasitic drugs in binary

aqueous-co-solvent mixtures.

Acknowledgments

We thank the Department of Pharmacy of the National University of Colombia for facilitating us the

equipment and laboratories used.

List of symbols

Latin letters

Cp Molar heat capacity

e11 Solvent-solvent interaction term

e13 Solvent-solute interaction term

e33 Solute-solute interaction term

R Gas constant

SMR Sulfamerazine

SMT Sulfamethazine

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T Absolute temperature

Tfus Absolute melting point

Thm Medium harmonic temperature

USP United States Pharmacopeia

x3 Solute mole fraction

V3 Molar volume of solute

w1 Mass fraction of ethanol (1) in the solvent binary mixture free of sulfamerazine (3) or

sulfamethazine (3).

Greek letters

3 Solute activity coefficient

δ Hildebrand solubility parameter

∆A→BG° Gibbs energy of transfer from solvent A to solvent B

∆A→BH° Enthalpy of transfer from solvent A to solvent B

∆A→BS° Entropy of transfer from solvent A to solvent B

∆fusH Enthalpy of fusion

∆mixG° Gibbs energy of mixing

∆mixH° Enthalpy of mixing

∆mixS° Entropy of mixing

∆solnG° Gibbs energy of solution

∆solnH° Enthalpy of solution

∆solnS° Entropy of solution

H Partial enthalpy contribution

TS Partial entropy contribution

�i Volume fraction of solvents in binary mixtures

φ1 Volume fraction of solvent in saturated the solution

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Table 1. Source and purities of the compounds used in this work.

Compound CAS FormulaMolar mass /

g mol–1 SourcePurity in mass

fractionSulfamerazine 127-79-7 C11H12N4O2S 264.31 Sigma Chemical Co. 0.990Sulfamethazine 57-68-1 C12H14N4O2S 278.33 Sigma Chemical Co. 0.990

Ethanol 64-17-5 C2H6O 46.07 Merck 0.998Water 7732-18-5 H2O 18.02 Obtained by distillation 1.000

Table 2. Experimental solubility of sulfamerazine (3) and sulfamethazine (3) in ethanol (1) + water (2) mixtures expressed in mole fraction including ideal solubility at several temperatures (± 0.05 K). Pressure = 73.9 ± 2.2 kPa.

Sulfamerazine (3)10,000 x3w1

a x1a

T = 293.15 K T = 298.15 K T = 303.15 K T = 308.15 K T = 313.15 K0.00 0.0000 0.134 ± 0.002 0.171 ± 0.003 0.209 ± 0.001 0.258 ± 0.003 0.316 ± 0.0040.10 0.0417 0.181 ± 0.002 0.232 ± 0.002 0.312 ± 0.007 0.405 ± 0.008 0.499 ± 0.0150.20 0.0891 0.343 ± 0.003 0.443 ± 0.004 0.578 ± 0.014 0.81 ± 0.03 1.00 ± 0.030.30 0.1436 0.752 ± 0.005 0.945 ± 0.021 1.243 ± 0.017 1.67 ± 0.03 2.07 ± 0.040.40 0.2068 1.405 ± 0.024 1.800 ± 0.003 2.35 ± 0.04 2.96 ± 0.04 3.64 ± 0.030.50 0.2812 2.414 ± 0.016 2.94 ± 0.05 3.64 ± 0.05 4.79 ± 0.01 5.86 ± 0.080.60 0.3698 3.335 ± 0.006 4.08 ± 0.06 5.02 ± 0.05 6.48 ± 0.03 7.67 ± 0.120.70 0.4772 4.134 ± 0.009 4.99 ± 0.07 5.92 ± 0.05 7.66 ± 0.04 9.27 ± 0.070.80 0.6101 4.43 ± 0.06 5.55 ± 0.07 6.45 ± 0.17 8.09 ± 0.18 9.43 ± 0.020.90 0.7788 3.82 ± 0.07 4.66 ± 0.03 5.59 ± 0.15 6.96 ± 0.04 8.11 ± 0.151.00 1.0000 2.74 ± 0.06 3.41 ± 0.10 4.09 ± 0.09 4.95 ± 0.04 6.36 ± 0.11

Ideal 46.2 ± 1.1 54.5 ± 1.3 64.1 ± 1.5 75.3 ± 1.8 88.1 ± 2.1Sulfamethazine (3)

10,000 x3w1 a x1

a

T = 293.15 K T = 298.15 K T = 303.15 K T = 308.15 K T = 313.15 K0.00 0.0000 0.222 ± 0.001 0.281 ± 0.001 0.367 ± 0.003 0.433 ± 0.001 0.555 ± 0.0040.10 0.0417 0.413 ± 0.006 0.495 ± 0.009 0.634 ± 0.010 0.836 ± 0.017 1.060 ± 0.0060.20 0.0891 0.769 ± 0.025 0.99 ± 0.03 1.20 ± 0.04 1.69 ± 0.07 2.21 ± 0.080.30 0.1436 1.566 ± 0.036 2.00 ± 0.06 2.68 ± 0.14 3.43 ± 0.09 4.66 ± 0.030.40 0.2068 2.814 ± 0.014 3.66 ± 0.07 4.84 ± 0.06 6.04 ± 0.08 8.25 ± 0.030.50 0.2812 4.80 ± 0.05 6.31 ± 0.04 7.79 ± 0.06 10.05 ± 0.18 13.19 ± 0.130.60 0.3698 7.27 ± 0.05 8.88 ± 0.11 11.45 ± 0.01 14.52 ± 0.09 18.83 ± 0.280.70 0.4772 9.49 ± 0.16 11.20 ± 0.07 14.69 ± 0.24 18.01 ± 0.09 22.93 ± 0.520.80 0.6101 10.89 ± 0.09 12.83 ± 0.14 16.62 ± 0.14 19.57 ± 0.04 24.72 ± 0.380.90 0.7788 10.02 ± 0.15 11.70 ± 0.29 14.16 ± 0.29 17.87 ± 0.23 21.47 ± 0.311.00 1.0000 7.47 ± 0.04 9.18 ± 0.07 10.87 ± 0.14 12.99 ± 0.23 16.23 ± 0.16

Ideal 88.6 ± 1.6 105.0 ± 1.9 124.1 ± 2.2 146.3 ± 2.6 172.0 ± 3.1a w1 and x1 are the mass and mole fractions of ethanol (1) in the solvent mixture free of sulfamerazine (3) or sulfamethazine (3), respectively.

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Table 3. Experimental solubility of sulfamerazine (3) and sulfamethazine (3) in ethanol (1) + water (2) mixtures expressed in molarity at several temperatures (± 0.05 K). Pressure = 73.9 ± 2.2 kPa.

Sulfamerazine (3)1,000 mol dm–3

w1 a x1

a

T = 293.15 K T = 298.15 K T = 303.15 K T = 308.15 K T = 313.15 K0.00 0.0000 0.742 ± 0.013 0.947 ± 0.015 1.152 ± 0.007 1.42 ± 0.02 1.74 ± 0.020.10 0.0417 0.928 ± 0.012 1.187 ± 0.012 1.59 ± 0.03 2.06 ± 0.04 2.54 ± 0.080.20 0.0891 1.62 ± 0.02 2.09 ± 0.02 2.72 ± 0.07 3.78 ± 0.16 4.67 ± 0.150.30 0.1436 3.26 ± 0.02 4.08 ± 0.09 5.34 ± 0.07 7.16 ± 0.12 8.82 ± 0.160.40 0.2068 5.51 ± 0.09 7.03 ± 0.01 9.18 ± 0.17 11.48 ± 0.15 14.07 ± 0.130.50 0.2812 8.51 ± 0.06 10.32 ± 0.17 12.72 ± 0.19 16.66 ± 0.02 20.27 ± 0.280.60 0.3698 10.46 ± 0.02 12.72 ± 0.18 15.58 ± 0.15 20.02 ± 0.10 23.59 ± 0.380.70 0.4772 11.40 ± 0.02 13.68 ± 0.20 16.18 ± 0.15 20.81 ± 0.12 25.03 ± 0.190.80 0.6101 10.63 ± 0.14 13.24 ± 0.18 15.31 ± 0.39 19.11 ± 0.42 22.17 ± 0.050.90 0.7788 7.83 ± 0.13 9.49 ± 0.06 11.34 ± 0.30 14.05 ± 0.09 16.28 ± 0.301.00 1.0000 4.70 ± 0.10 5.81 ± 0.17 6.94 ± 0.15 8.35 ± 0.07 10.66 ± 0.18

Sulfamethazine (3)1,000 mol dm–3

w1 a x1

a

T = 293.15 K T = 298.15 K T = 303.15 K T = 308.15 K T = 313.15 K0.00 0.0000 1.230 ± 0.007 1.557 ± 0.008 2.027 ± 0.017 2.390 ± 0.008 3.056 ± 0.0230.10 0.0417 2.11 ± 0.03 2.53 ± 0.04 3.23 ± 0.05 4.25 ± 0.08 5.38 ± 0.030.20 0.0891 3.63 ± 0.09 4.68 ± 0.16 5.62 ± 0.20 7.9 ± 0.3 10.3 ± 0.40.30 0.1436 6.77 ± 0.16 8.62 ± 0.26 11.5 ± 0.6 14.7 ± 0.4 19.85 ± 0.120.40 0.2068 11.04 ± 0.05 14.28 ± 0.28 18.80 ± 0.24 23.3 ± 0.3 31.76 ± 0.110.50 0.2812 16.88 ± 0.18 22.10 ± 0.14 27.11 ± 0.19 34.8 ± 0.6 45.5 ± 0.40.60 0.3698 22.74 ± 0.14 27.64 ± 0.34 35.43 ± 0.04 44.6 ± 0.3 57.4 ± 0.80.70 0.4772 26.1 ± 0.4 30.65 ± 0.18 40.0 ± 0.6 48.68 ± 0.24 61.6 ± 1.40.80 0.6101 26.07 ± 0.22 30.5 ± 0.3 39.3 ± 0.3 45.93 ± 0.09 57.6 ± 0.90.90 0.7788 20.5 ± 0.3 23.8 ± 0.6 28.7 ± 0.6 36.0 ± 0.5 42.9 ± 0.61.00 1.0000 12.97 ± 0.08 15.71 ± 0.12 18.44 ± 0.24 21.9 ± 0.4 27.14 ± 0.26a w1 and x1 are the mass and mole fractions of ethanol (1) in the solvent mixture free of sulfamerazine (3) or sulfamethazine (3), respectively.

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Table 4. Application of the Fedors’ method to estimate internal energy, molar volume, and Hildebrand solubility parameter of sulfamerazine (3) and sulfamethazine (3).

Sulfamerazine Sulfamethazine

GroupGroup number

E / kJ mol–1 V / cm3 mol–1 Group number

E / kJ mol–1 V / cm3 mol–1

–NH2 1 12.6 19.2 1 12.6 19.2–NH– 1 8.4 4.5 1 8.4 4.5=N– 2 2 x 11.7 2 x 5.0 2 2 x 11.7 2 x 5.0

–SO2– 1 25.6 19.5 1 25.6 19.5–CH3 1 4.71 33.5 2 2 x 4.71 2 x 33.5>C= 1 2 x 4.3 2 x –5.5 3 3 x 4.3 3 x –5.5

–CH= 3 2 x 4.3 2 x 13.5 1 4.3 13.5Phenylene ring 1 31.9 52.4 1 31.9 52.4Ring closure 1 1.05 16.0 1 1.05 16.0

Conjugate bond 3 3 x 1.67 3 x –2.2 3 3 x 1.67 3 x –2.2

Etotal = 129.91 Vtotal = 164.5 Etotal = 134.62 Vtotal = 179.0δtotal = (129,910/164.5)1/2 =

28.10 MPa1/2δtotal = (134,620/179.0)1/2 =

27.42 MPa1/2

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Table 5. Sulfamerazine (3) and sulfamethazine (3) activity coefficients (γ3) in ethanol (1) + water (2) cosolvente mixtures at several temperatures (± 0.05 K). Pressure = 73.9 ± 2.2 kPa.

a w1 and x1 are the mass and mole fractions of ethanol (1) in the solvent mixture free of sulfamerazine (3) or sulfamethazine (3), respectively.

Sulfamerazine (3)

w1 a x1

a T = 293.15 T = 298.15 T = 303.15 T = 308.15 T = 313.15

0.00 0.0000 345 ± 10 319 ± 9 308 ± 8 292 ± 8 279 ± 70.10 0.0417 255 ± 7 235 ± 6 206 ± 7 186 ± 6 177 ± 70.20 0.0891 135 ± 3 123 ± 3 111 ± 4 93 ± 5 88 ± 40.30 0.1436 61.5 ± 1.5 57.7 ± 1.9 51.6 ± 1.4 45.0 ± 1.3 42.6 ± 1.30.40 0.2068 32.9 ± 1.0 30.3 ± 0.7 27.3 ± 0.8 25.4 ± 0.7 24.2 ± 0.60.50 0.2812 19.2 ± 0.5 18.5 ± 0.5 17.6 ± 0.5 15.7 ± 0.4 15.0 ± 0.40.60 0.3698 13.9 ± 0.3 13.4 ± 0.4 12.8 ± 0.3 11.6 ± 0.3 11.5 ± 0.30.70 0.4772 11.2 ± 0.3 10.9 ± 0.3 10.8 ± 0.3 9.8 ± 0.2 9.5 ± 0.2

0.80 0.6101 10.4 ± 0.3 9.8 ± 0.3 9.9 ± 0.3 9.3 ± 0.3 9.3 ± 0.2

0.90 0.7788 12.1 ± 0.4 11.7 ± 0.3 11.5 ± 0.4 10.8 ± 0.3 10.9 ± 0.31.00 1.0000 16.9 ± 0.5 16.0 ± 0.6 15.7 ± 0.5 15.2 ± 0.4 13.8 ± 0.4

Sulfamethazine (3)

w1 a x1

a T = 293.15 T = 298.15 T = 303.15 T = 308.15 T = 313.15

0.00 0.0000 399 ± 8 373 ± 7 338 ± 7 338 ± 6 310 ± 60.10 0.0417 215 ± 5 212 ± 5 196 ± 5 175 ± 5 162 ± 30.20 0.0891 115 ± 6 106 ± 7 104 ± 4 86 ± 4 78 ± 30.30 0.1436 56.6 ± 1.7 52.5 ± 1.9 46.4 ± 2.5 42.7 ± 1.3 36.9 ± 0.70.40 0.2068 31.5 ± 0.6 28.7 ± 0.8 25.6 ± 0.6 24.2 ± 0.5 20.9 ± 0.40.50 0.2812 18.5 ± 0.4 16.6 ± 0.3 15.9 ± 0.3 14.6 ± 0.4 13.0 ± 0.30.60 0.3698 12.2 ± 0.2 11.8 ± 0.3 10.8 ± 0.2 10.1 ± 0.2 9.1 ± 0.20.70 0.4772 9.3 ± 0.2 9.4 ± 0.2 8.5 ± 0.2 8.1 ± 0.2 7.5 ± 0.2

0.80 0.6101 8.1 ± 0.2 8.2 ± 0.2 7.5 ± 0.1 7.5 ± 0.1 7.0 ± 0.2

0.90 0.7788 8.8 ± 0.2 9.0 ± 0.3 8.8 ± 0.2 8.2 ± 0.2 8.0 ± 0.21.00 1.0000 11.9 ± 0.2 11.4 ± 0.2 11.4 ± 0.3 11.3 ± 0.3 10.6 ± 0.2

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Table 6. Thermodynamic functions relative to solution process of sulfamerazine (3) and sulfamethazine(3) in ethanol (1) + water (2) co-solvent mixtures including ideal process at 303.0 ± 0.05 K. Pressure = 73.9 ± 2.2 kPa. .

Sulfamerazine (3)

w1 a x1

a ∆solnG° /kJ mol–1

∆solnH° /kJ mol–1

∆solnS° /J mol–1 K–1

T∆solnS° /kJ mol–1 H

b TS b

0.00 0.0000 27.2 ± 0.3 32.5 ± 0.4 17.5 ± 0.3 5.3 ± 0.1 0.859 0.1410.10 0.0417 26.2 ± 0.5 39.4 ± 0.7 43.6 ± 1.1 13.2 ± 0.3 0.749 0.2510.20 0.0891 24.5 ± 0.6 41.8 ± 1.0 56.8 ± 1.9 17.2 ± 0.6 0.708 0.2920.30 0.1436 22.6 ± 0.4 39.6 ± 0.7 55.9 ± 1.3 16.9 ± 0.4 0.700 0.3000.40 0.2068 21.1 ± 0.3 36.7 ± 0.5 51.5 ± 0.9 15.6 ± 0.3 0.702 0.2980.50 0.2812 19.9 ± 0.2 34.5 ± 0.7 48.2 ± 1.1 14.6 ± 0.3 0.702 0.2980.60 0.3698 19.1 ± 0.2 32.5 ± 0.6 44.2 ± 0.9 13.4 ± 0.3 0.708 0.2920.70 0.4772 18.6 ± 0.1 31.2 ± 0.7 41.4 ± 1.0 12.5 ± 0.3 0.713 0.2870.80 0.6101 18.5 ± 0.3 28.8 ± 0.6 34.2 ± 0.9 10.3 ± 0.3 0.736 0.2640.90 0.7788 18.9 ± 0.3 29.2 ± 0.5 34.1 ± 0.8 10.3 ± 0.2 0.739 0.2611.00 1.0000 19.6 ± 0.4 31.4 ± 0.8 38.8 ± 1.2 11.8 ± 0.4 0.727 0.273

Ideal 12.7 ± 0.3 24.6 ± 0.6 39.1 ± 1.3 11.9 ± 0.4 0.675 0.325Sulfamethazine (3)

w1 a x1

a ∆solnG° /kJ mol–1

∆solnH° /kJ mol–1

∆solnS° /J mol–1 K–1

T∆solnS° /kJ mol–1 H

b TS b

0.00 0.0000 25.8 ± 0.2 34.6 ± 0.6 28.9 ± 0.5 8.7 ± 0.2 0.798 0.2020.10 0.0417 24.3 ± 0.4 36.7 ± 0.9 41.1 ± 1.2 12.5 ± 0.4 0.747 0.2530.20 0.0891 22.6 ± 1.0 40.3 ± 1.6 58.6 ± 3.5 17.8 ± 1.1 0.694 0.3060.30 0.1436 20.7 ± 0.6 41.5 ± 1.0 68.4 ± 2.5 20.7 ± 0.7 0.667 0.3330.40 0.2068 19.3 ± 0.2 40.5 ± 0.7 70.0 ± 1.4 21.2 ± 0.4 0.656 0.3440.50 0.2812 18.0 ± 0.2 38.0 ± 0.6 65.9 ± 1.2 20.0 ± 0.4 0.655 0.3450.60 0.3698 17.0 ± 0.1 36.5 ± 0.6 64.3 ± 1.3 19.5 ± 0.4 0.652 0.3480.70 0.4772 16.5 ± 0.2 34.2 ± 0.8 58.4 ± 1.6 17.7 ± 0.5 0.659 0.3410.80 0.6101 16.2 ± 0.1 31.4 ± 0.7 50.3 ± 1.2 15.3 ± 0.4 0.673 0.3270.90 0.7788 16.5 ± 0.3 29.7 ± 0.8 43.7 ± 1.4 13.2 ± 0.4 0.692 0.3081.00 1.0000 17.2 ± 0.2 29.0 ± 0.5 38.9 ± 0.8 11.8 ± 0.3 0.711 0.289

Ideal 11.1 ± 0.2 25.3 ± 0.5 47.1 ± 1.2 14.3 ± 0.4 0.640 0.360a w1 and x1 are the mass and mole fractions of ethanol (1) in the solvent mixture free of sulfamerazine (3) or sulfamethazine (3), respectively.bH and TS are the relative contributions by enthalpy and entropy toward Gibbs energy of solution. These values were calculated by means of equations 7 and 8, respectively.

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Table 7. Thermodynamic functions of transfer of sulfamerazine (3) and sulfamethazine (3) from more polar solvents to less polar solvents in ethanol (1) + water (2) mixtures at 303.0 ± 0.05 K. Pressure = 73.9 ± 2.2 kPa. .

Sulfamerazine (3)w1

a x1a

A B A B∆A→BG° /kJ mol–1

∆A→BH° /kJ mol–1

∆A→BS° /J mol–1 K–1

T∆A→BS° /kJ mol–1

0.00 0.10 0.0000 0.0417 –1.0 ± 0.6 6.9 ± 0.8 26.1 ± 1.1 7.9 ± 0.3

0.10 0.20 0.0417 0.0891 –1.7 ± 0.8 2.4 ± 1.2 13.3 ± 2.2 4.0 ± 0.70.20 0.30 0.0891 0.1436 –1.9 ± 0.7 –2.2 ± 1.2 –0.9 ± 2.3 –0.3 ± 0.7

0.30 0.40 0.1436 0.2068 –1.5 ± 0.4 –2.9 ± 0.9 –4.4 ± 1.6 –1.3 ± 0.50.40 0.50 0.2068 0.2812 –1.2 ± 0.3 –2.2 ± 0.8 –3.3 ± 1.4 –1.0 ± 0.4

0.50 0.60 0.2812 0.3698 –0.8 ± 0.3 –2.0 ± 0.9 –4.1 ± 1.4 –1.2 ± 0.4

0.60 0.70 0.3698 0.4772 –0.5 ± 0.2 –1.3 ± 0.9 –2.8 ± 1.3 –0.8 ± 0.4

0.70 0.80 0.4772 0.6101 –0.2 ± 0.3 –2.4 ± 1.0 –7.3 ± 1.4 –2.2 ± 0.40.80 0.90 0.6101 0.7788 0.4 ± 0.4 0.4 ± 0.8 –0.1 ± 1.2 0.0 ± 0.4

0.90 1.00 0.7788 1.0000 0.8 ± 0.5 2.2 ± 0.9 4.8 ± 1.5 1.4 ± 0.4Sulfamethazine (3)

w1a x1

a

A B A B∆A→BG° /kJ mol–1

∆A→BH° /kJ mol–1

∆A→BS° /J mol–1 K–1

T∆A→BS° /kJ mol–1

0.00 0.10 0.0000 0.0417 –1.5 ± 0.4 2.2 ± 1.1 12.2 ± 1.3 3.7 ± 0.4

0.10 0.20 0.0417 0.0891 –1.7 ± 1.1 3.6 ± 1.9 17.5 ± 3.7 5.3 ± 1.1

0.20 0.30 0.0891 0.1436 –1.8 ± 1.2 1.1 ± 1.9 9.8 ± 4.3 3.0 ± 1.30.30 0.40 0.1436 0.2068 –1.5 ± 0.6 –1.0 ± 1.2 1.6 ± 2.8 0.5 ± 0.9

0.40 0.50 0.2068 0.2812 –1.3 ± 0.3 –2.5 ± 0.9 –4.1 ± 1.9 –1.2 ± 0.60.50 0.60 0.2812 0.3698 –0.9 ± 0.2 –1.4 ± 0.9 –1.6 ± 1.8 –0.5 ± 0.5

0.60 0.70 0.3698 0.4772 –0.6 ± 0.3 –2.4 ± 1.0 –5.9 ± 2.0 –1.8 ± 0.6

0.70 0.80 0.4772 0.6101 –0.3 ± 0.3 –2.7 ± 1.1 –8.1 ± 2.0 –2.4 ± 0.6

0.80 0.90 0.6101 0.7788 0.3 ± 0.3 –1.7 ± 1.1 –6.6 ± 1.9 –2.0 ± 0.60.90 1.00 0.7788 1.0000 0.7 ± 0.3 –0.8 ± 1.0 –4.8 ± 1.6 –1.5 ± 0.5a w1 and x1 are the mass and mole fractions of ethanol (1) in the solvent mixture free of sulfamerazine (3) or sulfamethazine (3), respectively; A and B are the more polar and less polar media, respectively.

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NH2

SNH

N

N

OO

R1

R2

Fig. 1. Molecular structure of the sulfonamides analyzed. Sulfadiazine (3): R1 and R2 = H. Sulfamerazine(3): R1 = H, R2 = CH3. Sulfamethazine (3): R1 and R2 = CH3.

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0.00

1.00

2.00

3.00

4.00

5.00

6.00

7.00

8.00

9.00

10.00

25.00 30.00 35.00 40.00 45.00 50.00

mix / MPa1/2

10,0

00 x

3

Fig. 2. Experimental solubility in mole fraction of sulfamerazine (x3) against the solubility parameter of the ethanol (1) + water (2) mixtures (δmix). (○) 293.15 K, (□) 298.15 K, (▲) 303.15 K, (●) 308.15 K, and (■) 313.15 K.

0.00

5.00

10.00

15.00

20.00

25.00

30.00

25.00 30.00 35.00 40.00 45.00 50.00

mix / MPa1/2

10,0

00 x

3

Fig. 3. Experimental solubility in mole fraction of sulfamethazine (x3) against the solubility parameter of the ethanol (1) + water (2) mixtures (δmix). (○) 293.15 K, (□) 298.15 K, (▲) 303.15 K, (●) 308.15 K, and (■) 313.15 K.

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4.0

6.0

8.0

10.0

12.0

14.0

16.0

18.0

0.00 0.10 0.20 0.30 0.40 0.50 0.60 0.70 0.80 0.90 1.00

w 1

m

ixG

° / k

J m

ol -

1

0.0

5.0

10.0

15.0

20.0

25.0

0.00 0.10 0.20 0.30 0.40 0.50 0.60 0.70 0.80 0.90 1.00

w 1

m

ixH

° / k

J m

ol -

1

-8.0

-6.0

-4.0

-2.0

0.0

2.0

4.0

6.0

8.0

10.0

0.00 0.10 0.20 0.30 0.40 0.50 0.60 0.70 0.80 0.90 1.00

w 1

mix

S°

/ kJ

mol

-1

Fig. 4. Thermodynamic quantities of mixing of sulfadiazine (▲), sulfamerazine (■), and sulfamethazine (●), in ethanol (1) + water (2) mixtures at 303.0 K as function of co-solvent mixtures composition.

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1.00

0.001.00 0.00

1.00

0.00

0.0

5.0

10.0

15.0

20.0

25.0

2.0 4.0 6.0 8.0 10.0 12.0 14.0 16.0 18.0

mixG ° / kJ mol -1

m

ixH

° / k

J m

ol -1

Fig. 5. ∆mixH° vs. ∆mixG° enthalpy-entropy compensation plot for dissolution process of sulfadiazine (3, ▲), sulfamerazine (3, ■), and sulfamethazine (3, ●), in ethanol (1) + water (2) co-solvent mixtures at 303.0 K.

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Solubility and solution thermodynamics of sulfamerazine and sulfamethazine in some ethanol +

water mixtures

Highlights

The solubility of sulfamerazine and sulfadiazine was maximal in a co-solvent mixture and minimal in

pure water.

Thermodynamic quantities of solution and mixing were obtained for sulfamerazine and sulfamethazine.

Non-linear plots of ∆mixH° vs. ∆mixG° compensation were found for sulfamerazine and sulfamethazine at

303.0 K.