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. 2023 Oct 22;28(20):7216. doi: 10.3390/molecules28207216

Thermodynamic Assessment of Triclocarban Dissolution Process in N-Methyl-2-pyrrolidone + Water Cosolvent Mixtures

Diego Ivan Caviedes-Rubio 1, Claudia Patricia Ortiz 2, Fleming Martinez 3, Daniel Ricardo Delgado 1,*
Editors: Piotr Cysewski, Tomasz Jeliński, Maciej Przybyłek
PMCID: PMC10609577  PMID: 37894697

Abstract

Solubility is one of the most important physicochemical properties due to its involvement in physiological (bioavailability), industrial (design) and environmental (biotoxicity) processes, and in this regard, cosolvency is one of the best strategies to increase the solubility of poorly soluble drugs in aqueous systems. Thus, the aim of this research is to thermodynamically evaluate the dissolution process of triclocarban (TCC) in cosolvent mixtures of {N-methyl-2-pyrrolidone (NMP) + water (W)} at seven temperatures (288.15, 293.15, 298.15, 303.15, 308.15, 313.15 and 318.15 K). Solubility is determined by UV/vis spectrophotometry using the flask-shaking method. The dissolution process of the TCC is endothermic and strongly dependent on the cosolvent composition, achieving the minimum solubility in pure water and the maximum solubility in NMP. The activity coefficient decreases from pure water to NMP, reaching values less than one, demonstrating the excellent positive cosolvent effect of NMP, which is corroborated by the negative values of the Gibbs energy of transfer. In general terms, the dissolution process is endothermic, and the increase in TCC solubility may be due to the affinity of TCC with NMP, in addition to the water de-structuring capacity of NMP generating a higher number of free water molecules.

Keywords: triclocarban, solubility, cosolvent, thermodynamics, N-methyl-2-pyrrolidone, water, modeling, simulation

1. Introduction

Triclocarban (TCC; 3,4,4-trichlorocarbanilide (Figure 1)) is a broad-spectrum antimicrobial agent, commonly used in personal care products, medical supplies, neonatal products, and even in civil infrastructure [1,2,3]. Despite its high effectiveness against Gram (+) and Gram (−) bacteria, TCC is considered a dangerous agent for public health by the Food and Drug Administration (FDA) and The European Commission (EC) because it is a potent endocrine disruptor [4,5].

Figure 1.

Figure 1

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Molecular structure of triclocarban.

In addition to being considered a dangerous agent for human health by the FDA and EC, TCC is listed in the NORMAN list [6] as an emerging contaminant of great danger to aquatic ecosystems, due to its recurrent presence in wastewater, sludge, and runoff [7,8,9].

Solubility is one of the most important physicochemical properties, which allows understanding the biopharmaceutical and pharmacokinetic processes of a drug, in addition to being related to design, formulation, preformulation, recrystallization, quantification, and quality evaluation processes [10,11,12]. On the other hand, in relation to solubility, the use of cosolvency is one of the most used techniques in the pharmaceutical industry to improve the solubility of drugs poorly soluble in aqueous systems [11,13,14]. Furthermore, from cosolvency studies, data of great relevance are determined, such as the dielectric requirement [15] and solubility parameter, as well as better understanding of possible molecular interactions through thermodynamic analysis [16,17] and preferential solvation [18,19,20,21,22].

In addition to interests in industrial processes, solubility also has relevance in environmental processes [23] since, from solubility, some important biological parameters can be determined, such as bioaccumulation [24].

Although TCC poses a risk to human health, some studies open the possibility of its use in the treatment of cancer and HIV [25,26], so the generation of physicochemical information regarding this drug is of great importance. Regarding the use of N-methyl-2-pyrrolidone (NMP) as a solvent, this solvent is used in the pharmaceutical industry, in the formulation of drugs for oral and transdermal administration, due to its great solubilizing power, stability, and miscibility with water in all proportions [27,28]. It is also used in extraction, purification, and crystallization processes of drugs [29,30]. Some interesting properties of NMP that make it an eco-friendly solvent are the possibility of being recycled by distillation and water extraction [29,31], in addition to being biodegradable [32] and biosynthesizable [33].

While some solubility data of TCC are reported in the literature, determined in pure solvents [34,35,36], cosolvent mixtures [37,38,39], aqueous systems [40] and in some solubilizing systems [41,42], the physicochemical information on the dissolution process of TCC is not complete, so the study of TCC solubility in {NMP (1) + W (2)} cosolvent mixtures will generate important information on issues of the solubility, cosolvency, and preferential solvation of TCC.

Therefore, the objective of the research is to evaluate the solubility of TCC in different {NMP (1) + W (2)} cosolvent mixtures at different temperatures, which is determined experimentally by UV/Vis spectrophotometry, and from the solubility data, the thermodynamic functions of solutions are calculated using the van ’t Hoff and Gibbs equation. Some of the most relevant results are that maximum solubility is achieved in a cosolvent mixture and that TCC is preferentially solvated by NMP in most cosolvent systems.

2. Results

2.1. Experimental Solubility (x3)

The experimental solubility data of TCC in cosolvent mixtures {NMP (1) + W (2)} are presented in Table 1 and Figure 2, where a strong dependence on the cosolvent composition can be observed. Thus, the minimum solubility is reached in pure water at 288.15 K and the maximum in NMP at 318.15 K.

Figure 2.

Figure 2

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Mole fraction of TCC (103 x3) depending on the mass fraction of NMP in the {NMP (1) + W (2)} mixtures free of TCC. •: 288.15 K; ∘: 293.15 K; △: 298.15 K; ▲: 303.15 K; ◻: 308.15 K; ◼: 313.15 K; 🟉: 318.15 K.

When analyzing the solubility behavior as a function of temperature, the solubility increases with the increase in temperature, indicating that the TCC solution process is endothermic. Regarding the solubility behavior of TCC as a function of the cosolvent composition, it increases as the solubility parameter of the mixture decreases by adding NMP (δ1 = 23.7 MPa1/2) [43]. Usually, the maximum solubility is reached when the solubility parameters of the drug and the solvent are equal; thus, according to the Fedors group contribution method, TCC has a solubility parameter of 26.5 MPa1/2 [36], so the maximum solubility of TCC should have been reached in a cosolvent mixture and not in pure NMP. However, Delgado et al. determined the TCC solubility parameter experimentally by studying the solubility of TCC in cosolvent mixtures {1,4-dioxane (1) + water (2)}, obtaining a solubility parameter for TCC of 21.92 MPa1/2 [44], so it is conjecturable that the maximum solubility of TCC in cosolvent mixtures {NMP (1) + W (2)} is reached in pure NMP since the solubility parameter of this solvent is greater than 21.92 MPa1/2.

Table 1.

Experimental solubility of TCC (3) in {NMP (1) + W (2)} cosolvent mixtures expressed in mole fraction at different temperatures (the values in parentheses are the standard deviations). Experimental pressure p: 0.096 MPa c.

w1 a Temperature/K b
288.15 293.15 298.15 303.15
0.00 1.96 × 10−9 d 2.38 × 10−9 d 2.85 × 10−9 d 3.78 × 10−9 d
0.05 4.088 (0.020) × 10−9 5.81 (0.07) × 10−9 7.52 (0.10) × 10−9 9.61 (0.10) × 10−9
0.10 8.71 (0.11) × 10−9 10.56 (0.09) × 10−9 15.48 (0.28) × 10−9 20.68 (0.2) × 10−9
0.15 1.66 (0.009) × 10−8 2.368 (0.013) × 10−8 3.143 (0.017) × 10−8 4.155 (0.042) × 10−8
0.20 3.82 (0.04) × 10−8 5.25 (0.04) × 10−8 7.09 (0.07) × 10−8 9.14 (0.04) × 10−8
0.25 8.5 (0.14) × 10−8 10.12 (0.13) × 10−8 14.751 (0.023) × 10−8 20.68 (0.06) × 10−8
0.30 1.669 (0.016) × 10−7 2.305 (0.026) × 10−7 3.32 (0.04) × 10−7 4.47 (0.04) × 10−7
0.35 4.424 (0.012) × 10−7 6.02 (0.06) × 10−7 7.99 (0.04) × 10−7 10.36 (0.04) × 10−7
0.40 0.996 (0.008) × 10−6 1.245 (0.007) × 10−6 1.783 (0.02) × 10−6 2.375 (0.018) × 10−6
0.45 2.363 (0.016) × 10−6 3.29 (0.04) × 10−6 4.27 (0.04) × 10−6 5.58 (0.04) × 10−6
0.50 5.76 (0.05) × 10−6 7.371 (0.023) × 10−6 9.64 (0.14) × 10−6 12.52 (0.05) × 10−6
0.55 1.258 (0.007) × 10−5 1.505 (0.01) × 10−5 2.205 (0.005) × 10−5 2.86 (0.029) × 10−5
0.60 2.827 (0.031) × 10−5 3.679 (0.032) × 10−5 4.884 (0.015) × 10−5 6.542 (0.035) × 10−5
0.65 6.89 (0.07) × 10−5 8.55 (0.09) × 10−5 10.64 (0.11) × 10−5 14.86 (0.05) × 10−5
0.70 1.459 (0.013) × 10−4 1.94 (0.021) × 10−4 2.628 (0.009) × 10−4 3.369 (0.03) × 10−4
0.75 3.456 (0.024) × 10−4 4.449 (0.033) × 10−4 5.73 (0.014) × 10−4 7.29 (0.04) × 10−4
0.80 7.48 (0.08) × 10−4 9.79 (0.1) × 10−4 12.28 (0.2) × 10−4 15.7 (0.13) × 10−4
0.85 1.631 (0.025) × 10−3 2.143 (0.009) × 10−3 2.728 (0.023) × 10−3 3.48 (0.018) × 10−3
0.90 3.63 (0.03) × 10−3 4.669 (0.03) × 10−3 6.004 (0.029) × 10−3 7.43 (0.13) × 10−3
0.95 8.01 (0.06) × 10−3 10.16 (0.09) × 10−3 13.09 (0.10) × 10−3 16.52 (0.10) × 10−3
1.00 1.741 (0.009) × 10−2 2.396 (0.017) × 10−2 2.777 (0.018) × 10−2 3.473 (0.015) × 10−2
w1 a Temperature/K b
308.15 313.15 318.15
0.00 5.72 × 10−9 d 7.48 × 10−9 d 9.28 × 10−9 d
0.05 12.39 (0.08) × 10−9 15.83 (0.1) × 10−9 23.42 (0.3) × 10−9
0.10 25.99 (0.27) × 10−9 33.2 (0.3) × 10−9 47.2 (0.8) × 10−9
0.15 5.48 (0.05) × 10−8 7.13 (0.07) × 10−8 9.25 (0.11) × 10−8
0.20 11.7 (0.12) × 10−8 16.87 (0.18) × 10−8 22.19 (0.25) × 10−8
0.25 26.48 (0.27) × 10−8 36.05 (0.37) × 10−8 46.7 (0.6) × 10−8
0.30 5.81 (0.08) × 10−7 7.68 (0.1) × 10−7 9.49 (0.07) × 10−7
0.35 15.35 (0.15) × 10−7 19.92 (0.19) × 10−7 23.3 (0.31) × 10−7
0.40 3.02 (0.02) × 10−6 3.93 (0.03) × 10−6 5.25 (0.01) × 10−6
0.45 7.028 (0.016) × 10−6 8.877 (0.02) × 10−6 13.18 (0.1) × 10−6
0.50 15.83 (0.11) × 10−6 22.96 (0.16) × 10−6 27.5 (0.28) × 10−6
0.55 3.55 (0.04) × 10−5 4.57 (0.05) × 10−5 5.979 (0.024) × 10−5
0.60 8.281 (0.025) × 10−5 10.415 (0.032) × 10−5 13.57 (0.2) × 10−5
0.65 18.82 (0.05) × 10−5 23.63 (0.07) × 10−5 31.47 (0.13) × 10−5
0.70 4.24 (0.04) × 10−4 5.186 (0.004) × 10−4 6.69 (0.05) × 10−4
0.75 9.176 (0.031) × 10−4 11.49 (0.04) × 10−4 15.4 (0.16) × 10−4
0.80 20.05 (0.17) × 10−4 25.07 (0.22) × 10−4 32.31 (0.32) × 10−4
0.85 4.388 (0.032) × 10−3 5.48 (0.04) × 10−3 6.79 (0.05) × 10−3
0.90 9.54 (0.06) × 10−3 11.89 (0.08) × 10−3 14.78 (0.11) × 10−3
0.95 20.33 (0.26) × 10−3 25.32 (0.33) × 10−3 31.84 (0.23) × 10−3
1.00 4.72 (0.04) × 10−2 5.585 (0.018) × 10−2 6.95 (0.11) × 10−2
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a w1 is the mass fraction of NMP (1) in the {NMP (1) + W (2)} mixtures free of TCC (3); b Standard uncertainty in temperature is u(T) = 0.05 K; c Standard uncertainty in pressure u(p) = 0.001 MPa; d Values taken from a reference [44].

Figure 2 shows the great cosolvent power of NMP, increasing the solubility of TCC by seven orders of magnitude from pure water to pure NMP. The low solubility of TCC in water may be due to the structuring of water around the non-polar groups of TCC [45]. When adding NMP, the solubility of TCC increases possibly due to two mechanisms. The first is the cosolvent effect of NMP, which in mixtures rich in water weakens the water structure, improving the solubility of TCC [46,47]; this effect is similar to the cosolvent action of ethanol, which is an excellent disruptor of the water structure. The second mechanism, which can occur in mixtures rich in NMD, is the possible formation of an NMP-TCC complex due to hydrophobic interactions [27,46,48] between non-polar groups of both NMP and TCC; NMD presents a relatively large and almost flat sector, which could enhance the formation of this possible NMP-TCC complex, which would theoretically favor the solubility of TCC.

A factor that can intervene in the change in solubility of a drug is the polymorphic changes or formation of solvates [49,50,51]; therefore, it is important to evaluate whether the change in the cosolvent composition promotes the formation of polymorphs.

A classic test to evaluate polymorphic changes is differential scanning calorimetry (DSC). In this vein, Figure 3 shows the DSC spectra of TCC from three solid phases in equilibrium with water, w0.5 and NMP and the commercial sample. Table 2 presents the temperature and enthalpy of the fusion results for each of the samples analyzed.

Figure 3.

Figure 3

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DSC thermograms of TCC.

Table 2.

The thermophysical properties of TCC obtained by the DSC.

Sample Enthalpy of Fusion,
ΔfusH/kJ·mol−1 		Melting Point,
Tfus/K		 Ref.
Original sample 41.3 ± 0.5 528.4 ± 0.5 This work
41.94 528.2 [52]
527.8 [39]
525 [44]
528.15–529.15 [53]
Water 41.9 ± 0.5 527.5 ± 0.5 This work
w0.50 42.2 ± 0.5 527.9 ± 0.5 This work
NMP 41.2 ± 0.5 528.6 ± 0.5 This work
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According to the temperature and enthalpy of the fusion results of the four samples, there is a relative deviation no greater than 1.2 %, so it is viable to assume that no polymorphic changes have occurred. In addition, the results agree with those reported in the literature and previous studies by the research group (Table 2).

2.2. Ideal Solubility and Activity Coefficients

In addition to evaluating deviations from ideality, from the activity coefficient (γ3), different molecular interactions that can occur in the solution process can also be assessed (solute–solute: e33, solute–solvent: e13 and solvent–solvent: e11) according to the equation proposed by Hildebrand and Wood Hildebrand and Wood (Equation (1)) [12,54,55]:

lnγ3=(e11+e33−2e13)V3ϕ12RT (1)

where V3 is the molar volume of the super-cooled liquid solute, and finally, ϕ1 is the volume fraction of the solvent, R is the gas constant, and T is the absolute temperature of the solution. As a first approximation, for relatively low solubilities (x3), the term V3ϕ12R−1T−1 may be considered constant; thus, γ3 depends mainly on e11, e33 and e13 [56]. The e11 and e33 terms are unfavorable for solubility, whereas the e13 term favors the solution process. Thus, the activity coefficient is calculated as the ratio between the ideal solubility, which depends exclusively on the physicochemical properties of the drug [57] and the experimental solubility (Equation (2)):

γ3=exp−ΔfHRTf−TTfT+ΔCpRTf−TT−ΔCpRlnTfTx3 (2)

where T and Tf are in K, ΔfH is the enthalpy of fusion (in kJ·mol−1) of the solute, R is the gas constant (in kJ·mol−1·K−1), and ΔCp is the differential heat capacity of fusion (in kJ K−1·mol−1) [57]. Some researchers like Hildebrand and Scott [58], Neau and Flynn [59], Neau et al. [60] and Opperhuizen et al. [61], assume ΔCp to be the entropy of fusion (ΔfS), which is calculated as ΔfH/Tm.

According to the results of γ3 (Table 3) the experimental solubility data of TCC in cosolvent mixtures {NMP (1) + water (2)} deviate strongly from ideality, reaching values up to 1.6 ×106 in pure water at 288.15 K. As the temperature increases from 288.15 to 318.15 K, γ3 decreases between 1.6 and 1.8 times possibly due to the increase in molecular agitation increasing the likelihood of particle collision, thus increasing the solubility of TCC and thereby decreasing γ3 [62,63]; when evaluating γ3 in terms of cosolvent composition, the increase in NMP in the cosolvent mixture produces a drastic decrease in γ3, up to 8.69 × 106 times in relation to the values in pure water, i.e., solute–solute (e33) and solvent–solvent (e11) interactions are stronger in more polar media and at the lowest study temperatures. As the polarity of the cosolvent system decreases as a result of the addition of NMP, solute–solvent molecular interactions (e13) increase, favoring the solubility of TCC reaching values of γ3 close to one (near-ideal behavior) between w1=0.75 and w1=0.85 from the mixture w1=0.85, and the values of γ3 are less than one, indicating a behavior that exceeds ideality, demonstrating the excellent cosolvent power of NMP.

Table 3.

Activity coefficient of TCC (3) in {NMP (1) + water (2)} cosolvent mixtures at different temperatures and pressure p = 0.096 MPa.

w1 a Temperature/K
288.15 293.15 298.15 303.15 308.15 313.15 318.15
0.00 1.56 × 106 1.52 × 106 1.49 × 106 1.32 × 106 1.02 × 106 9.08 × 105 8.51 × 105
0.05 7.50 × 105 6.22 × 105 5.65 × 105 5.18 × 105 4.70 × 105 4.29 × 105 3.37 × 105
0.10 3.52 × 105 3.42 × 105 2.74 × 105 2.41 × 105 2.24 × 105 2.05 × 105 1.67 × 105
0.15 1.85 × 105 1.53 × 105 1.35 × 105 1.20 × 105 1.06 × 105 9.52 × 104 8.53 × 104
0.20 8.02 × 104 6.89 × 104 5.99 × 104 5.44 × 104 4.97 × 104 4.02 × 104 3.56 × 104
0.25 3.61 × 104 3.57 × 104 2.88 × 104 2.41 × 104 2.20 × 104 1.88 × 104 1.69 × 104
0.30 1.84 × 104 1.57 × 104 1.28 × 104 1.11 × 104 1.00 × 104 8.84 × 103 8.32 × 103
0.35 6.93 × 103 6.00 × 103 5.32 × 103 4.81 × 103 3.79 × 103 3.41 × 103 3.39 × 103
0.40 3.08 × 103 2.90 × 103 2.38 × 103 2.10 × 103 1.92 × 103 1.73 × 103 1.51 × 103
0.45 1.30 × 103 1.10 × 103 9.94 × 102 8.93 × 102 8.28 × 102 7.65 × 102 5.99 × 102
0.50 5.33 × 102 4.90 × 102 4.41 × 102 3.98 × 102 3.68 × 102 2.96 × 102 2.87 × 102
0.55 2.44 × 102 2.40 × 102 1.93 × 102 1.74 × 102 1.64 × 102 1.49 × 102 1.32 × 102
0.60 1.08 × 102 98.2 87.0 76.1 70.3 65.2 58.2
0.65 44.5 42.3 39.9 33.5 30.9 28.7 25.1
0.70 21.0 18.6 16.2 14.8 13.7 13.1 11.8
0.75 8.87 8.12 7.41 6.83 6.34 5.91 5.13
0.80 4.10 3.69 3.46 3.17 2.90 2.71 2.44
0.85 1.88 1.69 1.56 1.43 1.33 1.24 1.16
0.90 0.845 0.774 0.707 0.67 0.61 0.571 0.534
0.95 0.383 0.356 0.325 0.301 0.286 0.268 0.248
1.00 0.176 0.151 0.153 0.143 0.123 0.122 0.114
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a w1 is the mass fraction of NMP (1) in the {NMP (1) + W (2)} mixtures free of TCC (3).

2.3. Thermodynamic Functions of Solution

From the experimental solubility data of TCC in cosolvent mixtures {NMP (1) + W (2)}, the enthalpy and Gibbs energy of solution are calculated using the van ’t Hoff–Krug equation (Equations (3) and (4)) [64,65,66,67,68]:

ΔsolnH°=−R∂lnx3∂T−1−Thm−1p (3)
ΔsolnG°=−RThm.intercept (4)

where Thm is the harmonic mean of the study temperatures (302.8 K) (calculated as Thm=n/∑i=nn1/T, where n is the number of temperatures studied), and the intercept is b=lnx3302.8 this value (b) is taken from the linear equation of the modified van ’t Hoff plot (lnx3=m·(T−1−Thm−1)+b), where “m” is the slope and “b” is the intercept.

From the values of ΔsolnH° and ΔsolnG°, ΔsolnS° is derived from the Gibbs equation as

ΔsolnS°=ΔsolnH°−ΔsolnG°Thm−1 (5)

From Equations (5) and (7), the contribution of the energetic and organizational components to the Gibbs energy is evaluated, and this contribution is corroborated through the Perlovich graphical method [69,70]:

ζH=|ΔsolnH°|(|TΔsolnS°|+|ΔsolnH°|)−1 (6)
ζTS=1−ζH (7)

According to the results, values greater than 0.5 for ζH indicate a greater contribution of the energy component to the solution process, that is, molecular interactions, such as indrogen bridges and van der Waals forces, represent a more relevant role than the effects related to the TΔsolnS° factor.

Table 4 shows the thermodynamic functions of the solution process of TCC (3) in {NMP (1) + W (2)} cosolvent mixtures. The Gibbs energy is positive in all cases and decreases from pure water to pure NMP. It is important to clarify that the standard Gibbs energy of solution is not an indicator of spontaneity; its positive value is a consequence of expressing solubility in molar fraction. Therefore, according to Equation (4), which technically is ΔsolnG°=−RThmlnx3, negative values are not obtained in any case. In relation to the uncertainties of the Gibbs energy, which is a propagation of the uncertainties of solubility and temperature, they are relatively small because the uncertainties of the solubility data values are also low.

Table 4.

Thermodynamic functions of the solution process of TCC (3) in {NMP (1) + W (2)} co-solvent mixtures at 302.18 K and pressure p = 0.096 MPa (the values in parentheses are the standard deviations).

w1 a ΔsolnG°/kJ·mol−1 ΔsolnH°/kJ·mol−1 ΔsolnS°/J·mol−1·K−1 TΔsolnS°/kJ/mol−1 ζH ζTS
0.00 48.6 b 41.55 b −23.33 b −7.06 b 0.85 b 0.15 b
0.05 46.5 (0.4) 42.1 (0.8) −14.37 (0.30) −4.35 (0.09) 0.91 0.09
0.10 44.7 (0.5) 42.9 (0.8) −5.85 (0.13) −1.77 (0.04) 0.96 0.04
0.15 42.8 (0.3) 43.1 (0.3) 0.962 (0.010) 0.291 (0.003) 0.99 0.01
0.20 40.8 (0.4) 44.2 (0.6) 11.13 (0.18) 3.37 (0.05) 0.93 0.07
0.25 38.9 (0.4) 44.8 (0.7) 19.6 (0.4) 5.94 (0.11) 0.88 0.12
0.30 37.0 (0.4) 44.6 (0.5) 25.3 (0.4) 7.67 (0.13) 0.85 0.15
0.35 34.6 (0.3) 43.8 (0.7) 30.2 (0.6) 9.13 (0.17) 0.83 0.17
0.40 32.7 (0.23) 42.5 (0.5) 32.4 (0.4) 9.82 (0.13) 0.81 0.19
0.45 30.49 (0.21) 41.6 (0.8) 36.6 (0.7) 11.08 (0.22) 0.79 0.21
0.50 28.4 (0.22) 40.6 (0.7) 40.3 (0.7) 12.20 (0.22) 0.77 0.23
0.55 26.45 (0.19) 40.1 (0.6) 45.2 (0.8) 13.70 (0.24) 0.75 0.25
0.60 24.36 (0.17) 39.85 (0.25) 51.2 (0.5) 15.49 (0.14) 0.72 0.28
0.65 22.27 (0.14) 39.0 (0.6) 55.1 (0.9) 16.69 (0.28) 0.7 0.3
0.70 20.24 (0.15) 38.2 (0.4) 59.4 (0.7) 17.99 (0.22) 0.68 0.32
0.75 18.2 (0.1) 37.3 (0.4) 63.0 (0.7) 19.09 (0.22) 0.66 0.34
0.80 16.27 (0.17) 36.8 (0.27) 67.8 (0.9) 20.53 (0.26) 0.64 0.36
0.85 14.3 (0.11) 36.1 (0.1) 72.1 (0.6) 21.85 (0.19) 0.62 0.38
0.90 12.34 (0.1) 35.7 (0.2) 77.0 (0.8) 23.31 (0.23) 0.6 0.4
0.95 10.39 (0.09) 34.9 (0.2) 81.0 (0.9) 24.52 (0.26) 0.59 0.41
1.00 8.39 (0.06) 34.7 (0.6) 87.0 (1.6) 26.3 (0.5) 0.57 0.43
Ideal 13.36 b 24.05 b 35.27 b 10.69 b 0.692 b 0.308 b
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a w1 is the mass fraction of NMP (1) in the {NMP (1) + W (2)} mixtures free of TCC (3); b Values taken from a reference [39].

As for ΔsolnH°, it is positive in all cases, indicating that the solution process is endothermic. Hence, as the temperature increases, the solubility of TCC increases. In pure water and in water-rich mixtures, ΔsolnS° is negative. This may be due to the structuring of water around the non-polar groups of TCC. From w1=0.15, ΔsolnS° takes positive values, indicating a possible destructuring of water by NMP. Therefore, the enthalpy increases from pure water to w1=0.25. This increase indicates the formation of bonds that may be water–water iterations (hydrophobic hydration). From w1=0.25, ΔsolnH° decreases, possibly due to the destructuring of water.

When evaluating the energetic and organizational contributions to Gibbs energy, the solution enthalpy is the major contributor, especially in water-rich systems, where its influence is greater than 90%. When analyzing the solution process through the Perlovich method (Figure 4), all values are recorded in sectors I (TΔsolnS°<ΔsolnH°) and VIII (ΔsolnH°>0 and TΔsolnS°<0, |TΔsolnS°|<|ΔsolnH°|), indicating that enthalpy contributes to the solution process to a greater extent.

Figure 4.

Figure 4

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Relation between enthalpy (ΔsolnH°) and entropy (ThmΔsolnS°) in terms of the process of TCC (3) solution in {NMP (1) + W (2)} cosolvent mixtures at 302.8 K. The isoenergetic curves for ΔsolnG° are represented by dotted lines.

2.4. Thermodynamic Functions of Transfer

The thermodynamic functions of the hypothetical transfer process of TCC from the medium of higher polarity to the medium of lower polarity (Table 5), are calculated as the difference between the value of the thermodynamic function (f is ΔtrG°, ΔtrH° or TΔtrS°) of the less polar medium and that of the more polar medium (Equation (8)):

Δtrf°=Δsolnflesspolar°−Δsolnfmorepolar° (8)

Table 5.

Thermodynamic functions of transfer of TCC (3) in {NMP (1) + W (2)} cosolvent mixtures at 302.8 K and pressure p = 0.096 MPa (the values in parentheses are the standard deviations).

More Polar (w1) → Less Polar (w1)a ΔtrG°/kJ·mol−1 ΔtrH°/kJ·mol−1 ΔtrS°/kJ·mol−1·K−1 TΔtrS°/kJ·mol−1
0.00 → 0.05 −2.2 (0.4) 0.6 (1.5) 8.9 (0.8) 2.71 (0.24)
0.05 → 0.10 −1.8 (0.7) 0.8 (1.1) 8.53 (0.33) 2.58 (0.10)
0.10 → 0.15 −1.8 (0.6) 0.2 (0.8) 6.81 (0.13) 2.06 (0.04)
0.15 → 0.20 −2.1 (0.5) 1.0 (0.6) 10.16 (0.18) 3.08 (0.05)
0.20 → 0.25 −1.9 (0.5) 0.7 (0.9) 8.5 (0.4) 2.58 (0.13)
0.25 → 0.30 −1.9 (0.6) −0.2 (0.9) 5.7 (0.6) 1.73 (0.17)
0.30 → 0.35 −2.3 (0.5) −0.9 (0.9) 4.8 (0.7) 1.46 (0.21)
0.35 → 0.40 −1.9 (0.4) −1.2 (0.9) 2.3 (0.7) 0.69 (0.21)
0.40 → 0.45 −2.21 (0.32) −1.0 (0.9) 4.1 (0.8) 1.25 (0.25)
0.45 → 0.50 −2.09 (0.31) −1.0 (1.0) 3.7 (1.0) 1.12 (0.31)
0.50 → 0.55 −1.95 (0.29) −0.5 (0.9) 5.0 (1.1) 1.5 (0.32)
0.55 → 0.60 −2.09 (0.25) −0.3 (0.7) 5.9 (0.9) 1.79 (0.28)
0.60 → 0.65 −2.09 (0.22) −0.9 (0.7) 4.0 (1.0) 1.2 (0.31)
0.65 → 0.70 −2.03 (0.20) −0.7 (0.7) 4.3 (1.2) 1.3 (0.4)
0.70 → 0.75 −2.03 (0.18) −0.9 (0.5) 3.6 (1.0) 1.1 (0.32)
0.75 → 0.80 −1.94 (0.20) −0.5 (0.5) 4.8 (1.1) 1.4 (0.3)
0.80 → 0.85 −1.97 (0.20) −0.65 (0.31) 4.3 (1.1) 1.32 (0.32)
0.85 → 0.90 −1.96 (0.15) −0.49 (0.24) 4.8 (1.0) 1.47 (0.3)
0.90 → 0.95 −1.95 (0.14) −0.75 (0.27) 4.0 (1.2) 1.2 (0.3)
0.95 → 1.00 −1.99 (0.11) −0.2 (0.6) 6.0 (1.9) 1.8 (0.6)
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a w1 is the mass fraction of NMP (1) in the {NMP (1) + W (2)} mixtures free of TCC (3).

By adding NMP to water, reducing the polarity of the medium, from w1=0.0 (pure water) to w1=0.25, the Gibbs energy of transfer is negative, indicating the preference of TCC for less polar media. This transfer process is favored by entropy (+) and disfavored by enthalpy (+). From w1=0.25 to w1=1.0 (pure NMP), the transfer process is also promoted to less polar media (ΔtrG° (−)). In addition, the transfer process is favored by enthalpy (−) and entropy (+). In general terms, the negative value of the Gibbs transfer energy in all cases reflects the positive cosolvent effect (increased solubility) of NMP.

When evaluating the contribution of ΔtrH° and TΔtrS° to the transfer process by the Perlovich graphical method (Figure 5), except for the process from w1=0.35 to w1=0.40 (Sector IV: ΔtrH°<0; TΔtrS°>0;|ΔtrH°|>|TΔtrS°|), the transfer process is driven by entropy (Sector II: ΔtrH°<TΔtrS° and Sector III: ΔtrH°<0; TΔtrS°>0; |TΔtrS°|>|ΔtrH°|). Although in the transfer (w1=0.35)→(w1=0.40) there is double favorability, at this point, the enthalpy of transfer contributes more to the process.

Figure 5.

Figure 5

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Relation between enthalpy (ΔtrH°) and entropy (ThmΔtrS°) of the process transfer of TCC (3) in {NMP (1) + W (2)} cosolvent mixtures at 302.8 K. The isoenergetic curves for ΔmixG° are represented by dotted lines.

2.5. Thermodynamic Functions of Mixing

The solution process involves two sub-processes. The first one consists of the melting of the solute remaining as a supercooled liquid and the formation of the cavity in the solvent to house the solute molecule. The second sub-process consists of the mixing of the liquids, solvent and supercooled solute (Figure 6).

Figure 6.

Figure 6

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Diagram of the hypothetical of solution process [71].

The solution process can be described by Equation (9):

ΔSolf°=Δmixf°+Δff302.8 (9)

Clearing Δmixf° of (9), we obtain

Δmixf°=Δsolnf°−Δff302.8 (10)

where f represents the Gibbs energy, enthalpy or entropy of mixing, and ff represents the thermodynamic functions of the fusion of TCC (3) and its cooling to the harmonic mean temperature, 302.8 K. As it has been described previously in the literature, in this research, the Δsolnf° values for the ideal solution processes were used instead of Δff302.8 [72].

The results of the thermodynamic functions of mixing are tabulated in Table 6. From pure water (w1=0.00) to w1=0.40, the mixing process discourages the solution process since the values of the enthalpy of mixing correspond to different types of interactions, one of which is the formation of the cavity, which is an endothermic process because energy must be supplied to break the solvent–solvent interactions (e11). This process disfavors the solubility of the solute. In pure water, and in mixtures rich in water, the mixing enthapy tends to decrease possibly due to the formation of water–water bonds, as a consequence of hydrophobic solvation around the non-polar groups of the solute, which agrees with the negative values of mixing entropy in aqueous and water-rich mixtures. From w1=0.40 to w1=0.85, the discouragement of the mixing process to the solution process persists; however, the values of ΔmixG° decrease due to the decrease in the enthalpy of mixing, possibly due to the increase in solute–solvent molecular interactions, which favor the solution process and, unlike the behavior in mixtures rich in water, an entropic favoring (+) is presented. Finally, from w1=0.90 to w1=1.00, the mixing process favors the solution process (ΔmixG° (−)); in this case, there is enthalpic discouragement and a greater entropic favoring.

Table 6.

Thermodynamic functions of mixing TCC (3) in {NMP (1) + W (2)} cosolvent mixtures at 302.8 K and pressure p = 0.096 MPa (the values in parentheses are the standard deviations).

w1 a ΔmixG°/kJ·mol−1 ΔmixH°/kJ·mol−1 ΔmixS°/kJ·mol−1K−1 TΔmixS°/kJ·mol−1
0.00 35.25 b 17.52 b −58.55 b −17.73 b
0.05 33.1 (0.4) 18.5 (0.8) −48.2 (0.3) −14.6 (0.1)
0.10 31.3 (0.5) 19.3 (0.8) −39.68 (0.18) −12.02 (0.06)
0.15 29.5 (0.3) 19.5 (0.3) −32.87 (0.14) −9.95 (0.04)
0.20 27.4 (0.4) 20.5 (0.6) −22.71 (0.22) −6.88 (0.07)
0.25 25.5 (0.4) 21.2 (0.7) −14.2 (0.4) −4.3 (0.12)
0.30 23.6 (0.4) 21 (0.5) −8.5 (0.4) −2.57 (0.13)
0.35 21.26 (0.27) 20.2 (0.7) −3.7 (0.6) −1.11 (0.18)
0.40 19.34 (0.24) 18.9 (0.5) −1.4 (0.4) −0.42 (0.14)
0.45 17.13 (0.22) 18 (0.8) 2.7 (0.7) 0.83 (0.22)
0.50 15.04 (0.23) 17.0 (0.7) 6.4 (0.7) 1.95 (0.22)
0.55 13.09 (0.19) 16.5 (0.6) 11.4 (0.8) 3.45 (0.24)
0.60 11.00 (0.18) 16.24 (0.25) 17.3 (0.5) 5.25 (0.15)
0.65 8.91 (0.15) 15.4 (0.6) 21.3 (0.9) 6.44 (0.28)
0.70 6.88 (0.16) 14.6 (0.4) 25.6 (0.8) 7.74 (0.23)
0.75 4.85 (0.12) 13.7 (0.4) 29.2 (0.7) 8.84 (0.23)
0.80 2.91 (0.18) 13.19 (0.27) 34.0 (0.9) 10.28 (0.26)
0.85 0.94 (0.12) 12.54 (0.15) 38.3 (0.7) 11.6 (0.2)
0.90 −1.02 (0.12) 12.05 (0.19) 43.2 (0.8) 13.07 (0.24)
0.95 −2.97 (0.11) 11.3 (0.19) 47.1 (0.9) 14.27 (0.26)
1.00 −4.97 (0.08) 11.1 (0.6) 53.1 (1.6) 16.1 (0.5)
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a w1 is the mass fraction of NMP (1) in the {NMP (1) + W (2)} mixtures free of TCC (3); b Values taken from a reference [39].

According to Perlovich’s analysis, from pure water to w1=0.85, the enthalpy of mixing governs the mixing process (Sector VIII: ΔmixH°>0, TΔmixS°<0, |ΔmixH°|>|TΔmixS°<0|, Sector I: ΔmixH°>TΔmixS°), and from w1=0.90 to w1=1.00, the entropy of mixing is the thermodynamic function that drives the mixing process (Sector II: ΔmixH°<TΔmixS°, |ΔmixH°|<|TΔmixS°|) (Figure 7).

Figure 7.

Figure 7

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Relation between enthalpy (ΔmixH°) and entropy (ThmΔmixS°) of the process mixing of TCC (3) in {NMP (1) + W(2)} cosolvent mixtures at 302.8 K. The isoenergetic curves for ΔmixG° are represented by dotted lines.

2.6. Enthalpy–Entropy Compensation Analysis

During the dissolution process, enthalpy changes occur that disfavor the process; however, these enthalpic changes are compensated by entropic changes as a result of non-covalent interactions between the solute and the solvents [73]. Sharp indicated that linear relations between the ΔsolnH° and TΔsolnS°, usually indicate strongly compensated processes [74].

The compensation between enthalpy and entropy generates defined trends through which the thermodynamic drive of the process can be identified [74,75].

According to Bustamante et al., the enthalpic–entropic compensation can be evaluated by plotting ΔsolnH° vs. ΔsolnG°. Thus, positive slopes indicate that the dissolution process is driven by the enthalpy of solution, and negative slopes indicate an entropic drive [64,76].

In this order of ideas, Figure 8 shows the behavior of the enthalpic–entropic compensation of the NMP solution process in {NMP (1) + W(2)} cosolvent mixtures. From pure water to w1=0.25, the process is driven by the solution entropy, and from w1=0.25 to pure NMP, the process is driven by the solution enthalpy.

Figure 8.

Figure 8

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Enthalpy–entropy compensation plot for the solubility of TCC (3) in {NMP (1) + W (2)} mixtures at Thm = 302.8 K.

3. Materials and Methods

3.1. Reagents

In this study, triclocarban (Sigma-Aldrich, Burlington, MA, USA; compound 3), N-methyl-2-pyrrolidone (Sigma-Aldrich, Burlington, MA, USA; the solvent component 1), double-distilled water (component 2) with conductivity lower than 2 μS cm−1, and ethanol (Sigma-Aldrich, Burlington, MA, USA) were used. Table 7 summarizes the sources and purities of the compounds studied.

Table 7.

Source and purities of the compounds used in this research.

Chemical Name CAS a Source Purity in Mass Fraction Analytic Technique b
Triclocarban 101-20-2 Sigma-Aldrich >0.990 HPLC
N-methyl-2-pyrrolidone 872-50-4 Sigma-Aldrich 0.998 GC
Water 25322-68-3
Ethanol 64-17-5 Sigma-Aldrich 0.998 GC
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a Chemical Abstracts Service Registry Number. b HPLC is high-performance liquid chromatography; GC is gas chromatography.

3.2. Preparation of Solvent Mixtures

All cosolvent mixtures {NMP (1) + W (2)} were prepared geometrically (mass fraction) using an analytical balance with a sensitivity of ±0.0001 g (RADWAG AS 220.R2, Krakow, Poland). In amber-colored bottles (capacity 10 mL), 19 mixtures were prepared, varying the mass fraction in 0.05. Three samples were prepared for each mixture.

3.3. Solubility Determination

Triclocarban solubility was determined according to the shake-flask method proposed by Higuchi and Connors [77,78,79]; the method is described in detail in some open access publications [80].

In general terms, the samples were saturated by adding an excess of TCC to ensure a liquid phase (saturated solution) and a solid phase (excessive drug). These were then deposited in a recirculation bath (Medingen K-22/T100, Medingen, Germany) at 288.15, 293.15, 298.15, 303.15, 308.15, 313.15 and 318.15 K for 72 h and periodically agitated. Subsequently, an aliquot of each sample was taken by a syringe, filtering the dispersion with a membrane with a pore diameter of 0.45 µm (Millipore Corp. Swinnex-13, Burlington, MA, USA), which was then diluted gravimetrically with ethanol to avoid the precipitation of TCC, followed by quantification by UV/Vis spectrophotometry (UV/Vis EMC-11- UV spectrophotometer, Duisburg, Germany) at 265 nm (wavelength of maximum absorbance) (see Supplementary Materials). Due to the low solubility of TCC in water-rich solvent mixtures (w1=0.05 to w1=0.40), the concentration was determined by the standard addition method. Thus, 10.00 g (m1) of a 10.00 μg/g TCC (C1) solution were taken, and 10.00 g of saturated TCC solution (m2) (unknown concentration) were added; the mixture (mf) was stirred to homogenize, and the absorbance of the final solution (mf) was determined. Then, the concentration was calculated using the calibration curve equation (Cf) (see Supplementary Materials). The concentration of the saturated solution (unknown concentration (Csoln-sat)) was determined by means of (Equation (11)):

Csoln-sat=Cf·mf−C1m1m2 (11)

3.4. Calorimetric Study

The enthalpy and melting temperature of four TCC samples were determined by differential scanning calorimetry (DSC 204 F1 Phoenix, Munich, Germany). The equipment was calibrated using Indium and Tin as standards, and an empty sealed pan was used as reference. A mass of approximately 10.0 mg of each sample was deposited in an aluminum crucible and placed in the calorimeter under a nitrogen flow of 10 mL min−1. The heating cycle was developed from 323 to 523 K, with a heating ramp of 10 K min−1. The solid samples in equilibrium with the saturated solution were dried at room temperature for 48 h under a continuous stream of dry air.

4. Conclusions

The solution process of triclocarban in {NMP (1) + W(2)} cosolvent mixtures is an endothermic process, heavily dependent on the cosolvent composition, demonstrating also the great solubilizing power of NMP, which increases the solubility of TCC up to seven orders of magnitude. This is corroborated by the thermodynamic transfer functions.

The solution process is endothermic, with enthalpic and entropic disfavor in water-rich mixtures and entropic favor in intermediate and NMP-rich mixtures, which can be corroborated with the mixing functions which disfavor the solution process in water-rich and intermediate mixtures and favor it in NMP-rich mixtures.

Finally, the solution process is strongly compensated and is driven by the entropy of the solution in water-rich mixtures, while in intermediate and NMP-rich mixtures, the solution process is driven by the solution enthalpy.

Acknowledgments

We thank the National Directorate of Research and National Committee for Research Development of the Universidad Cooperativa de Colombia for the financial support of the Project “Estudio termodinámico de la solubilidad de la isoniacida en sistemas acuosos polietilenglicol 200, 300, 400 y 600 e implementación de algoritmos de machine learning para la predicción de la solubilidad.” with code INV3171. We also thank the Universidad Cooperativa de Colombia, Sede Neiva, for facilitating the laboratories and equipment used.

Supplementary Materials

The following are available online at https://www.mdpi.com/article/10.3390/molecules28207216/s1, Figure S1: Calibration curve of triclocarban in absolute ethanol obtained at a wavelength of 265 nm. Table S1: Regression statistics. Table S2: ANOVA.

Click here for additional data file. (76.7KB, zip)

Author Contributions

Conceptualization, F.M. and D.R.D.; methodology, D.I.C.-R. and C.P.O.; software, D.I.C.-R.; validation, F.M. and D.R.D.; formal analysis, D.I.C.-R. and C.P.O.; investigation, D.I.C.-R. and C.P.O.; resources, D.R.D.; data curation, D.R.D.; writing—original draft preparation, F.M.; writing—review and editing, D.R.D. and F.M.; visualization, D.R.D.; supervision, F.M.; project administration, D.R.D.; funding acquisition, D.R.D. All authors have read and agreed to the published version of the manuscript.

Institutional Review Board Statement

Not applicable.

Informed Consent Statement

Not applicable.

Data Availability Statement

Data are contained within the article or Supplementary Materials.

Conflicts of Interest

The authors declare no conflict of interest.

Funding Statement

This research was funded by Universidad Cooperativa de Colombia grant number INV3171.

Footnotes

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