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Enthalpies and heat capacities of solution of methylpheophorbide, dioxidine and their conjugate in DMF at 298-318 K
Thermochimica Acta 669 (2018) 169–172
Contents lists available at ScienceDirect
Thermochimica Acta journal homepage: www.elsevier.com/locate/tca
Enthalpies and heat capacities of solution of methylpheophorbide, dioxidine and their conjugate in DMF at 298-318 K
T
Andrey V. Kustova,b, , Olga A. Antonovaa,b, Nataliya L. Smirnovaa,b, Irina S. Khudayevac, Dmitry V. Belykhc, Dmitry B. Berezinb ⁎
a
G.A. Krestov Institute of Solution Chemistry, Russian Academy of Sciences, Ivanovo, Russia Ivanovo State University of Chemistry and Technology, Ivanovo, Russia c Institute of Chemistry of Komi Republic Scientific Centre, Russian Academy of Sciences, Syktyvkar, Russia b
ARTICLE INFO
ABSTRACT
Keywords: N,N-Dimethylformamide Heterocycles Photosensitizers Enthalpy of solution Heat capacity of solution
This paper focuses on the results of the first accurate determination of enthalpies of solution of 9-ethenyl-14-ethyl-21(methoxycarbonyl)-3-{(O-methylcarbamoyl)ethyl} -4,8,13,18-tetramethyl-20-oxo-phorbine (Methylpheophorbide а, 1), [3-(hydroxymethyl)-1-oxido-4-oxoquinoxalin-4-ium-2-yl]methanol (Dioxidine, 2) and their conjugate 9-ethenyl14-ethyl-3-{[2-(3-hydroxymethyl)]-1-oxido-4-oxoquinoxalin-4-ium-2-methyleneoxy}carbamoyl}ethyl)-4,8,13,18-tetramethyl-20-oxo-phorbine (Conjugate 3) in liquid N,N-dimethylformamide (DMF) in the temperature range of 298318 K. Standard enthalpies and heat capacities of solution (ΔC°p) have been computed and compared with those for other solid and liquid non-electrolytes to highlight important features of solvation of potential photosensitizers in a simple model of a protein-like environment.
1. Introduction Photodynamic therapy (PDT) with an appropriate photosensitizer (PS) of a porphyrin, phthalocyanine or chlorophyll-type is a clinically approved therapeutic modality for treating many inflammatory and neoplastic processes [1–3]. Various skin diseases including acne and psoriasis, superficially located tumors, complicated wound infections caused by antibiotic resistant bacteria or fungi may be successfully treated with the PDT technique [4]. Most of the second-generation chlorophyll photosensitizers are appropriate derivatives of chlorin e6 or their metal complexes. These are usually synthesized from methylpheophorbide a or pheophorbide which are known to be an appropriate versatile synthetic base to develop various PSs [2]. [3-(hydroxymethyl)-1-oxido-4-oxoquinoxalin-4-ium-2-yl]methanol (Dioxidine) is a reserve quinoxaline N-oxide type antimicrobial drug which is important for treating multiresistent pathogenic microflora [5–7]. Dioxidine reveals broad-spectrum activity towards both anaerobic and mixed strains. However, due to its slight teratogenic and mutagenic action it is rarely used in therapeutic doses [5,6]. It is believed that covalent attachment of water-soluble dioxidine to hydrophobic methylpheophorbide a increases PS solubility in biological liquids and may enhance its dark toxicity towards different pathogens
both before and, especially, after the PDT treatment. Moreover, this conjugation should weaken side effects mentioned above. In this paper we describe synthesis of such a conjugate and study its thermochemical behavior in liquid DMF which is considered as an appropriate model for a protein-like environment [8]. These results are compared with those for dioxidine and methylpheophorbide a as well as with previously reported enthalpies of solution of blood porphyrins and non-electrolytes modeling their side-chains [9–15]. 2. Experimental 2.1. Chemicals DMF (Panreac, > 99 mass %) was dried with 4 Å molecular sieves for several days and then distilled under reduced pressure at 303 K, the middle fraction being selected. Karl Fisher titration showed that water content in the final product was lower than 0.02 mass %. The standard enthalpy of solution of DMF in water at 298 K was found to be equal to -15.25 kJ mol-1, which was in a fair agreement with literature values of -15.27 and -15.31 kJ mol-1 [16,17], respectively. Dioxidine 2 was obtained from its dosage form contained 1 mass % of the drug dissolved in water (Borisovskiy ZMP, Belarus). This solution
⁎ Corresponding author at: United Physical-Chemical Center of Solution, G.A. Krestov Institute of Solution Chemistry of Russian Academy of Sciences, 153045, Academicheskaya str., 1, Ivanovo, Russian Federation E-mail address: [email protected] (A.V. Kustov).
https://doi.org/10.1016/j.tca.2018.09.022 Received 6 April 2018; Received in revised form 17 September 2018; Accepted 28 September 2018 Available online 29 September 2018 0040-6031/ © 2018 Elsevier B.V. All rights reserved.
Thermochimica Acta 669 (2018) 169–172
A.V. Kustov et al.
Fig. 1. Molecular structures of methylpheophorbide a (1), dioxidine (2), and their conjugate (3).
was cooled down to 275 K. Then dioxidine crystals were separated from a solution using Schott filters, washed by cold water several times and dried under reduced pressure at 353 K for several days. Purity of the solute was checked with UV-, visible and 1H NMR (see the information given elsewhere [18]). 9-ethenyl-14-ethyl-3-{[2-(3-hydroxymethyl)]-1-oxido-4-oxoquinoxalin-4-ium-2-methyl-eneoxy}carbamoyl}ethyl)-4,8,13,18-tetramethyl20-oxo-phorbine (Conjugate 3) was obtained via chemical functionalization of methylpheophorbide а 1 [19] to attach the dioxidine 2 residue (see Fig. 1). The details of synthesis and purification of the pigment are presented in the Supplementary material file. The final product was identified with 1H NMR, MS-, IR- and UV-Vis spectra (see the Supplementary Material file). Both compounds 1, 3 were recrystallized from acetone and dried under reduced pressure at 353 K for several days. The description of all chemicals with their purity is given in Table 1. The structures of solutes 1-3 are illustrated in Fig. 1.
2.2. Apparatus and methods Enthalpies of solution were measured with the self-built isoperibol automated ampoule calorimeters and described several times before [9–15]. The calorimetric vessel was equipped with a calibration heater, a titanium stirrer and a thermistor. A glass ampoule containing a solute was attached to the stirrer and crushed against the vessel bottom to initiate the dissolution process. Thermistor resistance was directly measured by the Standard Temperature Measuring Instrument (BIC, Minsk). The detection limit of the apparatus was 10 μK. The temperature instability in the «Thermostat A 3» bath (BIC, Minsk) were less than 1 mK in the temperature range from 275 to 350 K. The enthalpies of solution were measured by a comparative method. An electrical calibration was carried out before each experiment. The calorimeter was tested by measuring enthalpies and heat capacities of solution of several non-electrolytes in water including urea, tetramethylurea and 1-
Table 1 Chemicals in this study. Chemical name
Abbreviation/ phase state
Source
Initial mass fraction
Purification method
Analysis method
Final mass fraction
Water 1-propanol
Liquid PrOH, liquid DMF, liquid
Aldrich
> 0.995
Bidistillation -
> 0.995
Panreac (Spain)
> 0.99
1, solid
(Chlorin, Russia)
∼0.98
Drying with 4 Å sieves, vacuum distillation Recrystallization, vacuum drying
As stated by the supplier As stated by the supplier 1
> 0.95
2, solid
Borisov. ZMP (Belarus)
0.01
3, solid
Synthesis
-
N,N-dimethylformamide 9-ethenyl-14-ethyl-21-(methoxycarbonyl)-3-{(Omethylcarbamoyl)ethyl} -4,8,13,18-tetramethyl-20-oxophorbine (Methylpheophorbide a) [3-(hydroxymethyl)-1-oxido-4-oxoquinoxalin-4-ium-2-yl] methanol (Dioxidine) 9-ethenyl-14-ethyl-3-{[2-(3-hydroxymethyl)]-1-oxido-4oxoquinoxalin-4-ium-2-methyleneoxy}carbamoyl}ethyl)4,8,13,18-tetramethyl-20-oxo-phorbine (Conjugate)
Recrystallization, washing, vacuum drying Column chromatography, recrystallization, vacuum drying
H NMR
> 0.99
As stated by the supplier
> 0.99
1
> 0.95
H NMR
Note. Purity of methylpheophorbide a of 98 % was provided by the supplier via the HQLC measurements. Our 1H NMR study indicates that purity of both macrocycles 1, 3 is of 95 % (see the Supplementary Material File). The value of 0.01 for dioxidine denotes an initial mass fraction of the solute in an ampoule with a sterile solution. 170
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Table 2 Experimental (ΔsolHm/kJ mol-1) and standard (ΔsolH°/kJ mol-1) enthalpies of solution of methylpheophorbide a, dioxidine and their conjugate in N,N-dimethylformamide at different temperatures. 298.15 K
308.15 K
318.15 K
mSolutea
ΔsolHm
mSolute
ΔsolHm
mSolute
ΔsolHm
Methylpheophorbide a 0.000301 0.000196 0.000201 0.000788 0.000691 0.000947 0.001620 0.001085 ΔsolH° = 18.34 ± 0.11b
18.12 18.31 18.60 18.31 18.54 18.31 18.34 18.18
0.000102 0.000100 0.000401 0.000426 0.000605
21.85 21.59 21.44 21.49 21.99
0.000300 0.000366 0.000485 0.000557
25.39 25.79 25.27 25.24
Dioxidine 0.013081 0.000423 0.005029 0.007274 0.004229 0.010419 0.002766 0.007540 ΔsolH° = 18.11 ± 0.10
18.23 18.34 17.99 17.97 18.09 18.19 18.10 17.97
0.001440 0.000298 0.000438 0.000278 0.000594 0.000362
Conjugate 0.000821 0.000252 0.000612 0.000785 ΔsolH° = -9.41 ± 0.12
-9.54 -9.44 -9.26 -9.40
a b
ΔsolH° = 21.67 ± 0.21 19.12 19.08 18.82 18.94 19.00 19.07
ΔsolH° = 19.01 ± 0.09 0.000294 0.000785 0.000685 0.000937 ΔsolH° = -6.16 ± 0.19
0.000472 0.000653 0.000716 0.001437 0.000854
20.03 19.74 20.18 20.08 19.87
ΔsolH° = 19.98 ± 0.16 -6.35 -5.94 -6.27 -6.07
0.000822 0.000390 0.000294 0.000937 ΔsolH° = -3.07 ± 0.19
-3.29 -2.85 -3.14 -2.98
From here on mSolute denotes solute molality, mol kg-1. Uncertainties for the experimental quantities represent the twice standard deviation of the mean.
propanol. Our results were found in a very good agreement with the recommended literature values (see the Supplementary Material File). Enthalpies of solution were measured in the range of solute molalities from m = 1⋅10-4 mol kg-1 to m = 1.3⋅10-2 mol kg-1. Enthalpies of solution were corrected to the side effects associated with an ampoule crushing on the vessel bottom and DMF vaporization into free space of the calorimetric vessel.
±2
ΔsolH° = 25.43 ± 0.25
(xmean
x i ) 2 / n (n
that the interaction of all species with DMF should be very similar, i.e. they interact with neighbor solvent molecules mainly via Van der Waals forces. Formation of H-bonds between hydroxyl group of the dioxidine residue and DMF also contributes to the total enthalpy change for compounds 2, 3. However, this effect is rather small and cannot lead to the negative enthalpies of solution for the conjugate because the ΔsolH° value is positive for pure dioxidine forming two hydrogen bonds with DMF molecules (see Fig. 1). Hence, we are able to draw a conclusion that a negative enthalpy change for compound 3 mainly arises from packing effects of conjugate molecules in a condensed state, i.e. the structure of a crystal lattice for this solute is more bulk than those for compounds 1, 2. The temperature dependences of the ΔsolH° values are well reproduced by the following linear equations:
1)
The experimental pressure is 101.33 kPa. Standard uncertainties for temperature, pressure and molality are u (T)= ± 0.01 K, u(p)= ± 0.5 kPa and u(mSolute)= ± 1·10-6mol kg-1, respectively.
ΔsolH°(1)/kJ mol-1 = 18.29(0.08) + 0.352(0.007)∙(T/K -298.15), (1) sf = 0.22 kJ mol-1
3. Results and Discussion The experimental ΔsolHm values at different temperatures and concentrations are given in Table 2. We see that experimental values do not depend on solute molality. This indicates that solute-solute correlations are too small to influence ΔsolHm values in this concentration range. Thus, standard enthalpies of solution or enthalpies of solution at an infinite dilution ΔsolH° can be simply computed by the usual way as the mean values in the range of experimental results. These quantities are also given in Table 2. We see that standard enthalpies of solution are large and positive for both compounds 1 and 2, the ΔsolH° values being nearly identical at the standard temperature. This solute behavior is very similar to that usually observed for non-electrolytes in aprotic solvents [20–22] indicating that the solute-solvent interactions in a solution do not compensate the energetic cost for disrupting solute-solute interactions in a condensed phase. In contrast, enthalpies of solution of compound 3 are negative for all the temperatures studied. The comparison of the solute structures illustrated in Fig. 1 indicates
ΔsolH°(2)/kJ mol-1 = 18.10(0.05) + 0.093(0.004)∙(T/K -298.15), sf = 0.14 kJ mol-1 (2) ΔsolH°(3)/ kJ mol-1= -9.38(0.08) + 0.317(0.006)∙(T/K -298.15), sf = 0.16 kJ mol-1, (3) where the first term from here on is the standard enthalpy of solution at the reference temperature of 298.15 K, the second one is the heat capacity of solution ΔC°р; values in brackets represent the standard deviation of the mean, sf is the standard deviation of the fit. Eqs. (1)–(3) show that the heat capacity of solution is positive for all cases. For small dioxidine molecule it is of 100 J mol-1 K-1, whereas for macrocycles 1, 3 the heat capacity is three times as larger. The comparison of Eqs. (1)–(3) and solute structures illustrated in Fig. 1 indicates that the heat capacity change do not correlate with the solute size because the ΔC°р values for methylpheophorbide a and its larger conjugate with dioxidine are nearly identical. This may arise from the fact that polar and 171
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apolar groups contribute in an opposite manner to the heat capacities of solution of chlorophyll PSs. Recently [9–15], we have studied the thermochemical behavior of four blood porphyrins and some non-electrolytes modeling their sidechains in liquid DMF and its mixtures with 1-octanol in a physiological temperature range. The temperature changes of the ΔsolH° values for some of these solutes are well represented by the following linear functions:
Appendix A. Supplementary data Supplementary material related to this article can be found, in the online version, at doi:https://doi.org/10.1016/j.tca.2018.09.022. References [1] P. Agostinis, K. Berg, K.A. Cengel, T.H. Foster, A.W. Girotti, S.O. Gollnick, S.T. Hahn, M.R. Hamblin, A. Juzeniene, D. Kessel, et al., Photodynamic therapy of cancer: an update, CA Cancer J Clin. 61 (2011) 250–281, https://doi.org/10.3322/ caac.20114. [2] A.S. Brandis, Y. Salomon, A. Schetz, Chlorophyll sensitizers in photodynamic therapy, in: B. Grimm, R.J. Porra, W. Rüdiger, H. Scheer (Eds.), Chlorophylls and bacteriochlorophylls: Biochemistry, biophysics, functions and application, Springer, Berlin, 2006, pp. 461–483. [3] L. Huang, T. Dai, M.R. Hamblin, Antimicrobial photodynamic inactivation and photodynamic therapy for infections, in: C.J. Gomer (Ed.), Photodynamic therapy. Methods and protocols, Springer, New York, Dordrecht Heidelberg, London, 2010, pp. 155–174. [4] M.A. Biel, Photodynamic therapy of bacterial and fungal biofilm infections, in: C.J. Gomer (Ed.), Photodynamic therapy. Methods and protocols, Springer, New York, Dordrecht Heidelberg, London, 2010, pp. 175–194. [5] G. Gupta, D.P. Singh, A.K. Mishra, Synthesis and antimicrobial activity of some quinoxaline derivative, Chem. Sci. Trans. 2 (2013) 1256–1261, https://doi.org/10. 7598/cst2013.507. [6] L.V. Bolshakov, Antibacterial activity of dioxidine under aero- and anaerobic conditions, Antibiot. Med. Biotekhnol. 31 (1986) 760–764. [7] L.P. Sycheva, M.A. Kovalenko, S.M. Sheremet’eva, A.D. Durnev, V.S. Zhurkov, Study of mutagenic activity of dioxidine by the polyorgan micronuclear method, Bull. Exp. Biol. Med. 138 (2004) 165–167, https://doi.org/10.1007/BF02694364. [8] M. Abbate, G. Barone, G. Castronuovo, P.J. Cheek, G. Giancola, T.E. Leslie, T.H. Lilley, Thermodynamic behaviour of some uncharged organic molecules in concentrated aqueous urea solutions and other polar solvents, Thermochim. Acta. 173 (1990) 261–272. [9] A.V. Kustov, N.L. Smirnova, M.B. Berezin, Standard enthalpies and heat capacities of ethyl acetate and deuteroporphyrin dimethylester solution in N,N-dimethylformamide at 298–318 K, Thermochim. Acta. 521 (2011) 224–226. [10] A.V. Kustov, N.L. Smirnova, O.A. Antonova, Enthalpies and heat capacities of ethyl acetate solution in water and in several organic solvents at 298–318 K, J. Solut. Chem. 41 (2012) 1008–1012. [11] A.V. Kustov, N.L. Smirnova, M.B. Berezin, Enthalpies and heat capacities of hematoporphyrin solutions in N,N-dimethylformamide and 1-octanol, Russ. J. Phys. Chem. A. 86 (2012) 895–897. [12] A.V. Kustov, N.L. Smirnova, M.B. Berezin, Preferable solvation of decane and benzene in 1-octanol – N,N-dimethylformamide mixed solvent, Russ. J. Phys. Chem. A. 88 (1) (2014) 57–61. [13] A.V. Kustov, N.L. Smirnova, D.B. Berezin, M.B. Berezin, Thermodynamics of solution of proto- and mezoporphyrins in N,N-dimethylformamide, J. Chem. Thermodyn. 89 (2015) 123–126. [14] A.V. Kustov, M.B. Berezin, Thermodynamics of solution of hemato- and deuteroporphyrins in N,N-dimethylformamide, J. Chem. Eng. Data. 58 (9) (2013) 2502–2505. [15] A.V. Kustov, N.L. Smirnova, Solvation of decane and benzene in mixtures of 1octanol and N,N-dimethylformamide, Russ. J. Phys. Chem. A. 90 (9) (2016) 1778–1781. [16] A. Rouw, G. Somsen, Solvation and hydrophobic hydration of alkylsubstituted ureas and amides in N,N-dimethylformamide + water mixtures, J. Chem. Soc., Faraday Trans. II 78 (1982) 3397–3408. [17] D.D. Macdonald, M.E. Estep, M.D. Smith, J.B. Hyne, Heats of solution and the influence of solutes on the temperature of maximum density of water, J. Solut. Chem. 3 (1974) 713–725. [18] A.V. Kustov, N.L. Smirnova, Thermodynamics of solution and partition of dioxidine in water and the water / 1-octanol biphasic system, J. Mol. Liquids. 248 (2017) 842–846. [19] A.V. Kustov, D.V. Belykh, N.L. Smirnova, I.S. Khudyaeva, D.B. Berezin, Partition of methylpheophorbide a, dioxidine and their conjugate in the 1-octanol/phosphate saline buffer biphasic system, J. Chem. Thermodyn. 115 (2017) 302–306. [20] M.A. Varfalomeev, I.T. Rakipov, B.N. Solomonov, Calorimetric investigation of hydrogen bonding of formamide and its methyl derivatives in organic solvents and water, Intern. J. Thermophys. 34 (2013) 710–724. [21] V.B. Novikov, D.I. Abaidullina, N.Z. Gainutdinova, M.A. Varfalomeev, B.N. Solomonov, Calorimetric determination of the enthalpy of specific interaction of chloroform with a number of proton-acceptor compounds, Russ. J. Phys. Chem. 80 (2006) 1790–1794. [22] B.N. Solomonov, V.B. Novikov, M.A. Varfolomeev, N.M. Mileshko, A new method for the extraction of specific interaction enthalpy from the enthalpy of solvation, J. Phys. Org. Chem. 18 (2005) 49–61.
-1
ΔsolH°(PDE)/kJ mol = 26.17(0.12)+ 0.558(0.009)∙(T/K -298.15), sf = 0.13 kJ mol-1 (4) ΔsolH°(HDEDE)/kJ mol-1 = 9.13(0.02)+ 0.199(0.001)∙(T/K-298.15), sf = 0.02 kJ mol-1 (5) ΔsolH°(EtOAc)/kJ mol-1 = 0.832(0.01)- 0.001(0.00003)∙(T/K-298.15), sf = 0.01 kJ∙mol-1 (6) ΔsolH°(Decane)/kJ mol-1 = -298.15), sf = 0.01 kJ mol-1
11.24(0.01)+
0.035(0.0002)∙(T/K (7)
ΔsolH°(Benzene)/kJ mol-1 = 0.11(0.00)+ 0.002(0.0001)∙(T/K -298.15), sf = 0.01 kJ mol-1, (8) where PDE and HDEDE are protoporphyrin dimethylester and hematoporphyrine dimethylether dimethylester, respectively. Eqs. (4)–(8) reveal two important features of the solute behavior which are worth noting. First, for large HDEDE containing four polar groups the heat capacity of solution is significantly smaller than for less polar PDE. The heat capacity of polar ethyl acetate is also smaller than the ΔC°р value for apolar decane. This clearly indicates that apolar groups of porphyrins increases the heat capacity of solution of, whereas polar ones contribute in an opposite manner. Second, for aromatic benzene the heat capacity of solution is much smaller than for aliphatic decane and approaches zero. Moreover, Eq. (8) shows that dissolution of benzene is accompanied by a nearly zero enthalpy change for the temperature range studied indicating that diluted benzene solutions in DMF are almost athermic. Thus, we see that both aromatic and, especially, aliphatic side-groups induce the heat capacity increase for porphyrin-type PSs. The appearance of polar ester groups leads to the significant decrease of heat capacity values. It is obvious, however, that the largest contribution to a positive heat capacity change arises from the macrocycle itself. It is quite clear that similar conclusions are valid for chlorophyll-type PSs. 4. Conclusions We have studied for the first time the thermochemical behavior of methylpheophorbide а, dioxidine and their conjugate in a simple model of a protein-like environment – liquid DMF in a physiological temperature range. These results are compared with the previously reported enthalpies and heat capacities of solution of blood porphyrins and some non-electrolytes modeling their side-chains. Heat capacities of solution of chlorophyll-type PSs and dioxidine are found to be positive but smaller than for porphyrins of a similar structure. Apolar sidechains give a positive contribution to the ΔC°р values, whereas polar functional groups contribute in an opposite manner. Large heat capacity values of macroheterocycles are not directly associated with the molecular size of the solute and rather with its solvophobicity. The more solvophobic the PS molecule, the more positive heat capacity of solution in DMF is. Acknowledgement This work was partially supported by the Russian Scientific Foundation (Grant 14-23-00204 ⊓)
172
Contents lists available at ScienceDirect
Thermochimica Acta journal homepage: www.elsevier.com/locate/tca
Enthalpies and heat capacities of solution of methylpheophorbide, dioxidine and their conjugate in DMF at 298-318 K
T
Andrey V. Kustova,b, , Olga A. Antonovaa,b, Nataliya L. Smirnovaa,b, Irina S. Khudayevac, Dmitry V. Belykhc, Dmitry B. Berezinb ⁎
a
G.A. Krestov Institute of Solution Chemistry, Russian Academy of Sciences, Ivanovo, Russia Ivanovo State University of Chemistry and Technology, Ivanovo, Russia c Institute of Chemistry of Komi Republic Scientific Centre, Russian Academy of Sciences, Syktyvkar, Russia b
ARTICLE INFO
ABSTRACT
Keywords: N,N-Dimethylformamide Heterocycles Photosensitizers Enthalpy of solution Heat capacity of solution
This paper focuses on the results of the first accurate determination of enthalpies of solution of 9-ethenyl-14-ethyl-21(methoxycarbonyl)-3-{(O-methylcarbamoyl)ethyl} -4,8,13,18-tetramethyl-20-oxo-phorbine (Methylpheophorbide а, 1), [3-(hydroxymethyl)-1-oxido-4-oxoquinoxalin-4-ium-2-yl]methanol (Dioxidine, 2) and their conjugate 9-ethenyl14-ethyl-3-{[2-(3-hydroxymethyl)]-1-oxido-4-oxoquinoxalin-4-ium-2-methyleneoxy}carbamoyl}ethyl)-4,8,13,18-tetramethyl-20-oxo-phorbine (Conjugate 3) in liquid N,N-dimethylformamide (DMF) in the temperature range of 298318 K. Standard enthalpies and heat capacities of solution (ΔC°p) have been computed and compared with those for other solid and liquid non-electrolytes to highlight important features of solvation of potential photosensitizers in a simple model of a protein-like environment.
1. Introduction Photodynamic therapy (PDT) with an appropriate photosensitizer (PS) of a porphyrin, phthalocyanine or chlorophyll-type is a clinically approved therapeutic modality for treating many inflammatory and neoplastic processes [1–3]. Various skin diseases including acne and psoriasis, superficially located tumors, complicated wound infections caused by antibiotic resistant bacteria or fungi may be successfully treated with the PDT technique [4]. Most of the second-generation chlorophyll photosensitizers are appropriate derivatives of chlorin e6 or their metal complexes. These are usually synthesized from methylpheophorbide a or pheophorbide which are known to be an appropriate versatile synthetic base to develop various PSs [2]. [3-(hydroxymethyl)-1-oxido-4-oxoquinoxalin-4-ium-2-yl]methanol (Dioxidine) is a reserve quinoxaline N-oxide type antimicrobial drug which is important for treating multiresistent pathogenic microflora [5–7]. Dioxidine reveals broad-spectrum activity towards both anaerobic and mixed strains. However, due to its slight teratogenic and mutagenic action it is rarely used in therapeutic doses [5,6]. It is believed that covalent attachment of water-soluble dioxidine to hydrophobic methylpheophorbide a increases PS solubility in biological liquids and may enhance its dark toxicity towards different pathogens
both before and, especially, after the PDT treatment. Moreover, this conjugation should weaken side effects mentioned above. In this paper we describe synthesis of such a conjugate and study its thermochemical behavior in liquid DMF which is considered as an appropriate model for a protein-like environment [8]. These results are compared with those for dioxidine and methylpheophorbide a as well as with previously reported enthalpies of solution of blood porphyrins and non-electrolytes modeling their side-chains [9–15]. 2. Experimental 2.1. Chemicals DMF (Panreac, > 99 mass %) was dried with 4 Å molecular sieves for several days and then distilled under reduced pressure at 303 K, the middle fraction being selected. Karl Fisher titration showed that water content in the final product was lower than 0.02 mass %. The standard enthalpy of solution of DMF in water at 298 K was found to be equal to -15.25 kJ mol-1, which was in a fair agreement with literature values of -15.27 and -15.31 kJ mol-1 [16,17], respectively. Dioxidine 2 was obtained from its dosage form contained 1 mass % of the drug dissolved in water (Borisovskiy ZMP, Belarus). This solution
⁎ Corresponding author at: United Physical-Chemical Center of Solution, G.A. Krestov Institute of Solution Chemistry of Russian Academy of Sciences, 153045, Academicheskaya str., 1, Ivanovo, Russian Federation E-mail address: [email protected] (A.V. Kustov).
https://doi.org/10.1016/j.tca.2018.09.022 Received 6 April 2018; Received in revised form 17 September 2018; Accepted 28 September 2018 Available online 29 September 2018 0040-6031/ © 2018 Elsevier B.V. All rights reserved.
Thermochimica Acta 669 (2018) 169–172
A.V. Kustov et al.
Fig. 1. Molecular structures of methylpheophorbide a (1), dioxidine (2), and their conjugate (3).
was cooled down to 275 K. Then dioxidine crystals were separated from a solution using Schott filters, washed by cold water several times and dried under reduced pressure at 353 K for several days. Purity of the solute was checked with UV-, visible and 1H NMR (see the information given elsewhere [18]). 9-ethenyl-14-ethyl-3-{[2-(3-hydroxymethyl)]-1-oxido-4-oxoquinoxalin-4-ium-2-methyl-eneoxy}carbamoyl}ethyl)-4,8,13,18-tetramethyl20-oxo-phorbine (Conjugate 3) was obtained via chemical functionalization of methylpheophorbide а 1 [19] to attach the dioxidine 2 residue (see Fig. 1). The details of synthesis and purification of the pigment are presented in the Supplementary material file. The final product was identified with 1H NMR, MS-, IR- and UV-Vis spectra (see the Supplementary Material file). Both compounds 1, 3 were recrystallized from acetone and dried under reduced pressure at 353 K for several days. The description of all chemicals with their purity is given in Table 1. The structures of solutes 1-3 are illustrated in Fig. 1.
2.2. Apparatus and methods Enthalpies of solution were measured with the self-built isoperibol automated ampoule calorimeters and described several times before [9–15]. The calorimetric vessel was equipped with a calibration heater, a titanium stirrer and a thermistor. A glass ampoule containing a solute was attached to the stirrer and crushed against the vessel bottom to initiate the dissolution process. Thermistor resistance was directly measured by the Standard Temperature Measuring Instrument (BIC, Minsk). The detection limit of the apparatus was 10 μK. The temperature instability in the «Thermostat A 3» bath (BIC, Minsk) were less than 1 mK in the temperature range from 275 to 350 K. The enthalpies of solution were measured by a comparative method. An electrical calibration was carried out before each experiment. The calorimeter was tested by measuring enthalpies and heat capacities of solution of several non-electrolytes in water including urea, tetramethylurea and 1-
Table 1 Chemicals in this study. Chemical name
Abbreviation/ phase state
Source
Initial mass fraction
Purification method
Analysis method
Final mass fraction
Water 1-propanol
Liquid PrOH, liquid DMF, liquid
Aldrich
> 0.995
Bidistillation -
> 0.995
Panreac (Spain)
> 0.99
1, solid
(Chlorin, Russia)
∼0.98
Drying with 4 Å sieves, vacuum distillation Recrystallization, vacuum drying
As stated by the supplier As stated by the supplier 1
> 0.95
2, solid
Borisov. ZMP (Belarus)
0.01
3, solid
Synthesis
-
N,N-dimethylformamide 9-ethenyl-14-ethyl-21-(methoxycarbonyl)-3-{(Omethylcarbamoyl)ethyl} -4,8,13,18-tetramethyl-20-oxophorbine (Methylpheophorbide a) [3-(hydroxymethyl)-1-oxido-4-oxoquinoxalin-4-ium-2-yl] methanol (Dioxidine) 9-ethenyl-14-ethyl-3-{[2-(3-hydroxymethyl)]-1-oxido-4oxoquinoxalin-4-ium-2-methyleneoxy}carbamoyl}ethyl)4,8,13,18-tetramethyl-20-oxo-phorbine (Conjugate)
Recrystallization, washing, vacuum drying Column chromatography, recrystallization, vacuum drying
H NMR
> 0.99
As stated by the supplier
> 0.99
1
> 0.95
H NMR
Note. Purity of methylpheophorbide a of 98 % was provided by the supplier via the HQLC measurements. Our 1H NMR study indicates that purity of both macrocycles 1, 3 is of 95 % (see the Supplementary Material File). The value of 0.01 for dioxidine denotes an initial mass fraction of the solute in an ampoule with a sterile solution. 170
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Table 2 Experimental (ΔsolHm/kJ mol-1) and standard (ΔsolH°/kJ mol-1) enthalpies of solution of methylpheophorbide a, dioxidine and their conjugate in N,N-dimethylformamide at different temperatures. 298.15 K
308.15 K
318.15 K
mSolutea
ΔsolHm
mSolute
ΔsolHm
mSolute
ΔsolHm
Methylpheophorbide a 0.000301 0.000196 0.000201 0.000788 0.000691 0.000947 0.001620 0.001085 ΔsolH° = 18.34 ± 0.11b
18.12 18.31 18.60 18.31 18.54 18.31 18.34 18.18
0.000102 0.000100 0.000401 0.000426 0.000605
21.85 21.59 21.44 21.49 21.99
0.000300 0.000366 0.000485 0.000557
25.39 25.79 25.27 25.24
Dioxidine 0.013081 0.000423 0.005029 0.007274 0.004229 0.010419 0.002766 0.007540 ΔsolH° = 18.11 ± 0.10
18.23 18.34 17.99 17.97 18.09 18.19 18.10 17.97
0.001440 0.000298 0.000438 0.000278 0.000594 0.000362
Conjugate 0.000821 0.000252 0.000612 0.000785 ΔsolH° = -9.41 ± 0.12
-9.54 -9.44 -9.26 -9.40
a b
ΔsolH° = 21.67 ± 0.21 19.12 19.08 18.82 18.94 19.00 19.07
ΔsolH° = 19.01 ± 0.09 0.000294 0.000785 0.000685 0.000937 ΔsolH° = -6.16 ± 0.19
0.000472 0.000653 0.000716 0.001437 0.000854
20.03 19.74 20.18 20.08 19.87
ΔsolH° = 19.98 ± 0.16 -6.35 -5.94 -6.27 -6.07
0.000822 0.000390 0.000294 0.000937 ΔsolH° = -3.07 ± 0.19
-3.29 -2.85 -3.14 -2.98
From here on mSolute denotes solute molality, mol kg-1. Uncertainties for the experimental quantities represent the twice standard deviation of the mean.
propanol. Our results were found in a very good agreement with the recommended literature values (see the Supplementary Material File). Enthalpies of solution were measured in the range of solute molalities from m = 1⋅10-4 mol kg-1 to m = 1.3⋅10-2 mol kg-1. Enthalpies of solution were corrected to the side effects associated with an ampoule crushing on the vessel bottom and DMF vaporization into free space of the calorimetric vessel.
±2
ΔsolH° = 25.43 ± 0.25
(xmean
x i ) 2 / n (n
that the interaction of all species with DMF should be very similar, i.e. they interact with neighbor solvent molecules mainly via Van der Waals forces. Formation of H-bonds between hydroxyl group of the dioxidine residue and DMF also contributes to the total enthalpy change for compounds 2, 3. However, this effect is rather small and cannot lead to the negative enthalpies of solution for the conjugate because the ΔsolH° value is positive for pure dioxidine forming two hydrogen bonds with DMF molecules (see Fig. 1). Hence, we are able to draw a conclusion that a negative enthalpy change for compound 3 mainly arises from packing effects of conjugate molecules in a condensed state, i.e. the structure of a crystal lattice for this solute is more bulk than those for compounds 1, 2. The temperature dependences of the ΔsolH° values are well reproduced by the following linear equations:
1)
The experimental pressure is 101.33 kPa. Standard uncertainties for temperature, pressure and molality are u (T)= ± 0.01 K, u(p)= ± 0.5 kPa and u(mSolute)= ± 1·10-6mol kg-1, respectively.
ΔsolH°(1)/kJ mol-1 = 18.29(0.08) + 0.352(0.007)∙(T/K -298.15), (1) sf = 0.22 kJ mol-1
3. Results and Discussion The experimental ΔsolHm values at different temperatures and concentrations are given in Table 2. We see that experimental values do not depend on solute molality. This indicates that solute-solute correlations are too small to influence ΔsolHm values in this concentration range. Thus, standard enthalpies of solution or enthalpies of solution at an infinite dilution ΔsolH° can be simply computed by the usual way as the mean values in the range of experimental results. These quantities are also given in Table 2. We see that standard enthalpies of solution are large and positive for both compounds 1 and 2, the ΔsolH° values being nearly identical at the standard temperature. This solute behavior is very similar to that usually observed for non-electrolytes in aprotic solvents [20–22] indicating that the solute-solvent interactions in a solution do not compensate the energetic cost for disrupting solute-solute interactions in a condensed phase. In contrast, enthalpies of solution of compound 3 are negative for all the temperatures studied. The comparison of the solute structures illustrated in Fig. 1 indicates
ΔsolH°(2)/kJ mol-1 = 18.10(0.05) + 0.093(0.004)∙(T/K -298.15), sf = 0.14 kJ mol-1 (2) ΔsolH°(3)/ kJ mol-1= -9.38(0.08) + 0.317(0.006)∙(T/K -298.15), sf = 0.16 kJ mol-1, (3) where the first term from here on is the standard enthalpy of solution at the reference temperature of 298.15 K, the second one is the heat capacity of solution ΔC°р; values in brackets represent the standard deviation of the mean, sf is the standard deviation of the fit. Eqs. (1)–(3) show that the heat capacity of solution is positive for all cases. For small dioxidine molecule it is of 100 J mol-1 K-1, whereas for macrocycles 1, 3 the heat capacity is three times as larger. The comparison of Eqs. (1)–(3) and solute structures illustrated in Fig. 1 indicates that the heat capacity change do not correlate with the solute size because the ΔC°р values for methylpheophorbide a and its larger conjugate with dioxidine are nearly identical. This may arise from the fact that polar and 171
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apolar groups contribute in an opposite manner to the heat capacities of solution of chlorophyll PSs. Recently [9–15], we have studied the thermochemical behavior of four blood porphyrins and some non-electrolytes modeling their sidechains in liquid DMF and its mixtures with 1-octanol in a physiological temperature range. The temperature changes of the ΔsolH° values for some of these solutes are well represented by the following linear functions:
Appendix A. Supplementary data Supplementary material related to this article can be found, in the online version, at doi:https://doi.org/10.1016/j.tca.2018.09.022. References [1] P. Agostinis, K. Berg, K.A. Cengel, T.H. Foster, A.W. Girotti, S.O. Gollnick, S.T. Hahn, M.R. Hamblin, A. Juzeniene, D. Kessel, et al., Photodynamic therapy of cancer: an update, CA Cancer J Clin. 61 (2011) 250–281, https://doi.org/10.3322/ caac.20114. [2] A.S. Brandis, Y. Salomon, A. Schetz, Chlorophyll sensitizers in photodynamic therapy, in: B. Grimm, R.J. Porra, W. Rüdiger, H. Scheer (Eds.), Chlorophylls and bacteriochlorophylls: Biochemistry, biophysics, functions and application, Springer, Berlin, 2006, pp. 461–483. [3] L. Huang, T. Dai, M.R. Hamblin, Antimicrobial photodynamic inactivation and photodynamic therapy for infections, in: C.J. Gomer (Ed.), Photodynamic therapy. Methods and protocols, Springer, New York, Dordrecht Heidelberg, London, 2010, pp. 155–174. [4] M.A. Biel, Photodynamic therapy of bacterial and fungal biofilm infections, in: C.J. Gomer (Ed.), Photodynamic therapy. Methods and protocols, Springer, New York, Dordrecht Heidelberg, London, 2010, pp. 175–194. [5] G. Gupta, D.P. Singh, A.K. Mishra, Synthesis and antimicrobial activity of some quinoxaline derivative, Chem. Sci. Trans. 2 (2013) 1256–1261, https://doi.org/10. 7598/cst2013.507. [6] L.V. Bolshakov, Antibacterial activity of dioxidine under aero- and anaerobic conditions, Antibiot. Med. Biotekhnol. 31 (1986) 760–764. [7] L.P. Sycheva, M.A. Kovalenko, S.M. Sheremet’eva, A.D. Durnev, V.S. Zhurkov, Study of mutagenic activity of dioxidine by the polyorgan micronuclear method, Bull. Exp. Biol. Med. 138 (2004) 165–167, https://doi.org/10.1007/BF02694364. [8] M. Abbate, G. Barone, G. Castronuovo, P.J. Cheek, G. Giancola, T.E. Leslie, T.H. Lilley, Thermodynamic behaviour of some uncharged organic molecules in concentrated aqueous urea solutions and other polar solvents, Thermochim. Acta. 173 (1990) 261–272. [9] A.V. Kustov, N.L. Smirnova, M.B. Berezin, Standard enthalpies and heat capacities of ethyl acetate and deuteroporphyrin dimethylester solution in N,N-dimethylformamide at 298–318 K, Thermochim. Acta. 521 (2011) 224–226. [10] A.V. Kustov, N.L. Smirnova, O.A. Antonova, Enthalpies and heat capacities of ethyl acetate solution in water and in several organic solvents at 298–318 K, J. Solut. Chem. 41 (2012) 1008–1012. [11] A.V. Kustov, N.L. Smirnova, M.B. Berezin, Enthalpies and heat capacities of hematoporphyrin solutions in N,N-dimethylformamide and 1-octanol, Russ. J. Phys. Chem. A. 86 (2012) 895–897. [12] A.V. Kustov, N.L. Smirnova, M.B. Berezin, Preferable solvation of decane and benzene in 1-octanol – N,N-dimethylformamide mixed solvent, Russ. J. Phys. Chem. A. 88 (1) (2014) 57–61. [13] A.V. Kustov, N.L. Smirnova, D.B. Berezin, M.B. Berezin, Thermodynamics of solution of proto- and mezoporphyrins in N,N-dimethylformamide, J. Chem. Thermodyn. 89 (2015) 123–126. [14] A.V. Kustov, M.B. Berezin, Thermodynamics of solution of hemato- and deuteroporphyrins in N,N-dimethylformamide, J. Chem. Eng. Data. 58 (9) (2013) 2502–2505. [15] A.V. Kustov, N.L. Smirnova, Solvation of decane and benzene in mixtures of 1octanol and N,N-dimethylformamide, Russ. J. Phys. Chem. A. 90 (9) (2016) 1778–1781. [16] A. Rouw, G. Somsen, Solvation and hydrophobic hydration of alkylsubstituted ureas and amides in N,N-dimethylformamide + water mixtures, J. Chem. Soc., Faraday Trans. II 78 (1982) 3397–3408. [17] D.D. Macdonald, M.E. Estep, M.D. Smith, J.B. Hyne, Heats of solution and the influence of solutes on the temperature of maximum density of water, J. Solut. Chem. 3 (1974) 713–725. [18] A.V. Kustov, N.L. Smirnova, Thermodynamics of solution and partition of dioxidine in water and the water / 1-octanol biphasic system, J. Mol. Liquids. 248 (2017) 842–846. [19] A.V. Kustov, D.V. Belykh, N.L. Smirnova, I.S. Khudyaeva, D.B. Berezin, Partition of methylpheophorbide a, dioxidine and their conjugate in the 1-octanol/phosphate saline buffer biphasic system, J. Chem. Thermodyn. 115 (2017) 302–306. [20] M.A. Varfalomeev, I.T. Rakipov, B.N. Solomonov, Calorimetric investigation of hydrogen bonding of formamide and its methyl derivatives in organic solvents and water, Intern. J. Thermophys. 34 (2013) 710–724. [21] V.B. Novikov, D.I. Abaidullina, N.Z. Gainutdinova, M.A. Varfalomeev, B.N. Solomonov, Calorimetric determination of the enthalpy of specific interaction of chloroform with a number of proton-acceptor compounds, Russ. J. Phys. Chem. 80 (2006) 1790–1794. [22] B.N. Solomonov, V.B. Novikov, M.A. Varfolomeev, N.M. Mileshko, A new method for the extraction of specific interaction enthalpy from the enthalpy of solvation, J. Phys. Org. Chem. 18 (2005) 49–61.
-1
ΔsolH°(PDE)/kJ mol = 26.17(0.12)+ 0.558(0.009)∙(T/K -298.15), sf = 0.13 kJ mol-1 (4) ΔsolH°(HDEDE)/kJ mol-1 = 9.13(0.02)+ 0.199(0.001)∙(T/K-298.15), sf = 0.02 kJ mol-1 (5) ΔsolH°(EtOAc)/kJ mol-1 = 0.832(0.01)- 0.001(0.00003)∙(T/K-298.15), sf = 0.01 kJ∙mol-1 (6) ΔsolH°(Decane)/kJ mol-1 = -298.15), sf = 0.01 kJ mol-1
11.24(0.01)+
0.035(0.0002)∙(T/K (7)
ΔsolH°(Benzene)/kJ mol-1 = 0.11(0.00)+ 0.002(0.0001)∙(T/K -298.15), sf = 0.01 kJ mol-1, (8) where PDE and HDEDE are protoporphyrin dimethylester and hematoporphyrine dimethylether dimethylester, respectively. Eqs. (4)–(8) reveal two important features of the solute behavior which are worth noting. First, for large HDEDE containing four polar groups the heat capacity of solution is significantly smaller than for less polar PDE. The heat capacity of polar ethyl acetate is also smaller than the ΔC°р value for apolar decane. This clearly indicates that apolar groups of porphyrins increases the heat capacity of solution of, whereas polar ones contribute in an opposite manner. Second, for aromatic benzene the heat capacity of solution is much smaller than for aliphatic decane and approaches zero. Moreover, Eq. (8) shows that dissolution of benzene is accompanied by a nearly zero enthalpy change for the temperature range studied indicating that diluted benzene solutions in DMF are almost athermic. Thus, we see that both aromatic and, especially, aliphatic side-groups induce the heat capacity increase for porphyrin-type PSs. The appearance of polar ester groups leads to the significant decrease of heat capacity values. It is obvious, however, that the largest contribution to a positive heat capacity change arises from the macrocycle itself. It is quite clear that similar conclusions are valid for chlorophyll-type PSs. 4. Conclusions We have studied for the first time the thermochemical behavior of methylpheophorbide а, dioxidine and their conjugate in a simple model of a protein-like environment – liquid DMF in a physiological temperature range. These results are compared with the previously reported enthalpies and heat capacities of solution of blood porphyrins and some non-electrolytes modeling their side-chains. Heat capacities of solution of chlorophyll-type PSs and dioxidine are found to be positive but smaller than for porphyrins of a similar structure. Apolar sidechains give a positive contribution to the ΔC°р values, whereas polar functional groups contribute in an opposite manner. Large heat capacity values of macroheterocycles are not directly associated with the molecular size of the solute and rather with its solvophobicity. The more solvophobic the PS molecule, the more positive heat capacity of solution in DMF is. Acknowledgement This work was partially supported by the Russian Scientific Foundation (Grant 14-23-00204 ⊓)
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