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Solvatochromic and preferential solvation of fluorescein in some water-alcoholic mixed solvents
Journal of Molecular Liquids 190 (2014) 126–132
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Solvatochromic and preferential solvation of fluorescein in some water-alcoholic mixed solvents Fereshteh Naderi a,⁎, Ali Farajtabar b, Farrokh Gharib c a b c
Department of Chemistry, Shahr-e Qods branch, Islamic Azad University, Tehran, Iran Department of Chemistry, Jouybar branch, Islamic Azad University, Jouybar, Iran Department of Chemistry, Faculty of Sciences, Shahid Beheshti University, Tehran, Evin, Iran
a r t i c l e
i n f o
Article history: Received 27 August 2013 Received in revised form 17 October 2013 Accepted 31 October 2013 Available online 14 November 2013 Keywords: Fluorescein Solvatochromism Preferential solvation Hydrogen bonding
a b s t r a c t It is known that the hydrogen bonding has a significant effect on the spectral properties of fluorescein, however its relevance to solution behavior and preferential solvation of fluorescein is unclear in mixed solvents. In this work, absorption spectra of fluorescein dianion have been investigated in pure water, methanol, ethanol, 1-propanol, 2-propanol, dimethylsulfoxide, N,N-dimethylformamide, acetonitrile and in binary mixtures of water with the mentioned alcohols at 298.15 K. Spectral changes were interpreted in terms of specific and nonspecific solute-solvent interactions. The energy of electronic transition in maximum absorption (ET) was calculated in each binary mixture. The extent and importance of each of solute–solvent interactions to ET were analyzed in the framework of the linear solvation energy relationships. Preferential solvation was detected as a nonideal behavior of ET curve respective to the analytical mole fraction of alcohols in all binary mixtures. The ET values were fitted to a solvent exchange model to calculate the preferential solvation parameters. The preference of fluorescein dianion to be solvated by one of the solvating species relative to others was discussed in terms of solvent–solvent and solute–solvent interactions. © 2013 Elsevier B.V. All rights reserved.
1. Introduction Study on the solvation process opens the way to further understand the solvent effects on the kinetics and thermodynamics of reactions. The behavior of solute in solution is a function of all types of interactions taken place within the solvation shell. Evidently, solvation in mixed solvent is more complicated than pure solvents because of the interference of the solvent–solvent interactions and diversity of solute–solvent interactions. Interaction between solvent molecules can create new solvating complexes whose properties and their structures are distinct and different from those in the pure solvent. Furthermore, the molecules of solute can usually interact to different degrees with each solvent component. At this case, the composition of solvation shell near close to solute, sometimes called local composition, can differ from that of the bulk composition of mixed solvent. Therefore a specific feature of solvation in the mixture of solvents emerges as a result of the phenomenon which is termed as preferential solvation [1,2]. In other words, preferential solvation arises when the solute interacts with more molecules of one solvent component in its cybotactic region than the others in comparison of the bulk composition. According to the above explanation, the comprehension of preferential solvation is a key issue in solvation processes in mixed solvents.
⁎ Corresponding author. Tel./fax: +98 21 46842938. E-mail address: [email protected] (F. Naderi). 0167-7322/$ – see front matter © 2013 Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.molliq.2013.10.028
Solvatochromism, which describes spectral changes of solute induced by changes in solvent polarity, provides a convenient way for the monitoring of interactions that occur in the cybotactic region of solute [2–5]. In particular, the electronic transition energy of the solute in the absorption maxima named as polarity scale, ET, comprises all sorts of interactions between solute and solvent components including specific (such as hydrogen bonding and electron donor acceptor interactions) and nonspecific (polarity–polarizability effects) interactions [2,4,6–8]. The shifts of the absorption maxima with changing the solvent composition and ET values can therefore reveal the nature of the solute–solvent as well as solvent–solvent interactions present in the mixed solvents. Accordingly, ET parameters of a solute can be manipulated in order to elucidate the preferential solvation phenomena in the mixtures. Various theoretical and experimental models to the problem of preferential solvation have been developed in the literature. Between theoretical approaches, the Kirkwood–Buff theory, the quasi lattice-quasi chemical theory developed by Marcus, the dielectric enrichment, the competitive preferential solvation theory of Nagy, and the stepwise solvent exchange model of Covington has been the basis of many papers [9–14]. The solvent exchange model has been proposed by Bosch and Roses to investigate preferential solvation experimentally [15–17]. This model is an extension of the stepwise solvent exchange model of Connors which derives equations that relate ET values of a solute with the solvent composition [18]. In this model, competition between different solvent species to solvate the solute is described by some exchange
F. Naderi et al. / Journal of Molecular Liquids 190 (2014) 126–132
equilibria in vicinity of solute. An equilibrium constant, named preferential solvation parameter, is defined for each exchange reaction, which relates the mole fraction of solvents in solvation shell to that in the bulk mixture. In addition, formation of solvating complexes is postulated from solvent–solvent interaction on the microsphere of solvation. Recently, El-Seoud has modified this model by considering the formation of solvating complexes explicitly in both cybotactic region and bulk mixture, since to subtle description of solvation, all of species in the solvation shell should be considered in equilibrium with the same species in the bulk mixture [19–21]. Fluorescein, a xanthene dye, is a popular fluorescent probe with widespread applications to various areas such as biology, chemistry and technology [22–26]. Fluorescein and fluorescein derivatives have intense absorbance and fluorescence characteristics with high quantum yield in the visible region, low toxicity and the ability of attaching to some biomolecular and ionic species. The presence of seven prototropic forms with different photophysical properties has been identified for fluorescein depending on the pH of the solution [27–29]. Predominant form of fluorescein under neutral and alkaline solutions is fluorescein dianion (FD) which has high fluorescence quantum yield and large absorption coefficient [29]. FD, Fig. 1, is commercially available in sodium salt, known as uranine or resorcinolphthalein sodium. The spectral characteristics of FD are important factors which made it useful in water tracing, laser tuning, ophthalmic photography and fluorescein angiography [30–32]. Accordingly solvatochromic properties of fluorescein have been subject of some studies in aqueous and nonaqueous solvents [33–38]. As reported in literature, position and intensity of absorbance and fluorescence emission maxima in FD are generally dependent on the hydrogen bonding propensity of the solvent. To achieve a further insight into the solution behavior of FD, in this work solvatochromic properties as well as preferential solvation of FD were studied in some pure solvents and binary mixtures at 298.15 K. 2. Experimental 2.1. Materials and methods All pure solvents including methanol, ethanol, 1-propanol, 2propanol, dimethylsulfoxide (DMSO), N,N-dimethylformamide (DMF) and acetonitrile (ACN) were supplied as analytical reagent grade from Merck. Fluorescein disodium and potassium hydroxide were obtained from Sigma. NaOH solution was prepared from a titrisol solution (Merck). All dilute solutions were prepared from double-distilled water with a conductance equal to (1.3 ± 0.1) μS. Spectrophotometric measurements were performed on a UV–vis Shimadzu 2100 spectrophotometer with a Pentium 4 computer and using thermostated matched 10 mm quartz cells at 298.15 K. All aqueous binary mixtures were carefully prepared by weighing at the required molar ratio with an electrical balance accurate to ±0.1 mg. Alkaline solution of solvents was prepared by solving suitable amount of KOH into 50 mL of pure and binary mixed solvents. Stock solution of FD was prepared in ethanol. 50 μL of the stock solution was transferred to -
O
O
127
5 mL glass volumetric tubes. After the evaporation of the ethanol under reduced pressure, 3.5 mL of the pure or the binary solvent was pipetted into the tubes and the mixtures were sonicated to a clear and homogenous solution and then kept in the dark. The final concentration of the solute in the tubes was obtained 6.14 μM. To exclude oxygen from the system, a stream of purified nitrogen gas was passed through solutions prior to each spectrophotometric measurement. The electronic absorption spectra of FD were recorded over the wavelength range of (400–600) nm at a rate of 140 nm min−1 with a slit width of 2 nm. At least three replicate spectrophotometric measurements were done for each solution with an accuracy of ±0.05 nm. 3. Results and discussion 3.1. Solvatochromism of FD in pure and aqueous binary solvents Absorption spectra of FD in pure water, methanol, ethanol, 1propanol and 2-propanol are shown in Fig. 2. In water solution, spectrum of FD exhibits a maximum band at 490.8 nm accompanied with a shoulder located around 470 nm, which assigned to the electronic transition from the ground state (S0) to the first (S1) and second excited singlet state (S2) respectively. The predominant electronic transition is mainly described by the HOMO → LUMO excitation, with a small contribution from HOMO-5 → LUMO [27]. The structures of ground and excited states of FD have been provided in the literature [27,39,40]. These results have indicated that both HOMO and LUMO are localized on the xanthene moiety, and the first excited state (S1), with largest oscillator strength, corresponds mainly to the electron transition from HOMO → LUMO. On going from water to 2-propanol, the position of absorption maximum shifts toward longer wavelengths, with an increase in the absorbance intensity. The spectrum of FD in aqueous solutions of methanol, ethanol, 1-propanol and 2-propanol was also investigated over the entire range of the mole fraction of alcoholic solvent. A regular bathochromic shift was observed as the mole fraction of alcohol increases in the aqueous solution for all binary aqueous mixtures in this work. Quantum chemical calculations have provided evidence that the dipole moment of FD increases upon electronic transition from the ground to excited state in gas phase; the excited state is more polar than the ground state [41]. The polarity effect rules that the more polar excited state is a more stable state in media of high polarity; the reverse is true for the less polar ground state. Therefore, in an opposite direction of what was observed for FD in this work, one could reasonably expect that a hypsochromic shift should emerge as the polarity of media decreases upon a solvent change from water to alcoholic media. This suggests that the nonspecific interactions have no pronounced
O
COO-
Fig. 1. The molecular structure of fluorescein dianion.
Fig. 2. Absorption spectra of FD (6.14 μM) in water, methanol, ethanol, 1-propanol and 2-propanol at 298.15 K.
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media increases due to approximately pure electrostatic interactions. This contradictory is attributed to the difference over the nature of the ground and excited state; the ground state has appreciably a more hydrogen band accepting capacity, while the excited state is a more polar state. It seems that the hydrogen bonding interaction is also responsible for the reduction in the absorption intensity on going from alcoholic media to water, because hydrogen bonding in xanthene ring reduces the resonance of the aromatics cycles. To quantitative representation, the wavelength of the maximum absorbance, λmax, in a particular solvent mixture was calculated from the first derivative of the spectrum data and then transformed into Eq. (1) to calculate the molar electronic transition energies, ET values.
effect on the spectral shift of FD. On the other hand, the FD has several hydrogen bond acceptor sites on its molecular structure, which provide opportunities for the specific solute–solvent interactions. Literature, in consistent with our observation, shows that the electronic spectra of FD experience a blue shift when the hydrogen bond donating power of milieu increases [33–38]. This can be explained by this fact that the acidity of fluorescein in the excited state is significantly greater than that in the ground state, implying that the basicity of FD decreases upon excitation [28,29,42]. Moreover upon photoexcitation, electron densities at the oxygen atoms on the xanthene ring (tricycle) decrease in the excited state. Accordingly, the hydrogen bond accepting nature of FD in excited state is considerably lower than that corresponding to ground sate; thus, as expected, the absorption spectra shift to the shorter wavelength upon hydrogen bonding formation, due to more stabilization of ground state related to the excited state when the hydrogen bond donating power of media increases. The above controversy reveals that the contribution of nonspecific electrostatic solute–solvent interaction to the spectral shift of FD decreases dramatically in the presence of the specific solute–solvent interactions; so the effect of the solvent polarity appears to be masked by the strong hydrogen bonding formation between the FD and the solvent in protic solvents. To gain further verification, the spectral shift of FD was examined in pure aprotic solvents including DMSO, DMF and ACN, in which no significant hydrogen bonding exists between the FD and the solvent, due to the lack of any hydrogen bond donor sites on the FD; seemingly the nonspecific electrostatic interactions appear to be predominant in the absence of the hydrogen bonding interactions. Maximum in the absorption was observed at 515, 519 and 521 nm for ACN, DMF and DMSO respectively, where the polarity of solvent increases as the order ACN b DMF b DMSO. In agreement with the above explanation, a bathochromic shift is clearly seen when the polarity of medium increases, because the dipole–dipole interaction makes the more polar excited state favorably more stable than the ground state. Therefore, in aprotic solvents where hydrogen bond donation is excluded, the spectral shift of FD can be explained in terms of a change in the polarity of the medium. Overall, the net results of hydrogen bonding and electrostatic interactions explain the spectral behavior of FD in solvent. In protic solvents, the hydrogen bonding interactions are dominant for the solute; the main absorption band shifts hypsochromically with increasing the hydrogen bond donating power of media, whereas in aprotic solvents, the absorption bond undergoes a bathochromic shift as the polarity of
−1 ¼ hcNA =λmax ðnmÞ ¼ 28591:5=λmax ðnmÞ ET kcal·mol
ð1Þ
where h, c and NA are Planck's constant, the velocity of light and Avogadro's number, respectively. The values of the ET as a function of mole fraction of alcoholic solvent are presented in Table 1 and Fig. 3. Solvent effect was treated within the framework of the linear solvation energy relationships (LSER) established by Kamlet, Abboud and Taft, in which each of specific and nonspecific interactions have a linear contribution to the total solvation energy of solvent dependent phenomena [43]. In nonchloro-substituted aliphatic solvents, LSER takes the form of Eq. (2). ET ¼ A0 þ aα þ bβ þ pπ
ð2Þ
The α, β and π* are the Kamlet and Taft solvatochromic parameters (KAT) which have been developed for scaling the hydrogen-bond donor acidity, hydrogen-bond acceptor basicity and dipolarity/polarizability of solvents respectively [44–46]. The A0, a, b, and p are regression coefficients; A0 is intercept whereas a, b and p quantify the sensitivity of ET values to the acidity, basicity and dipolarity/polarizability of solvent respectively. The number of parameters in Eq. (2) depends physically on the nature of solute and solvent and the significance of the solute– solvent interactions. The KAT parameters for binary mixtures were extracted from the literature and inserted in Table 1 [47]; the value of α, and π* for 1-propanol is 0.84, 0.90 and 0.54 respectively [2]; no literature data was found for aqueous solutions of 1-propanol. The ET values of FD for each aqueous solution were correlated with solvent properties by
Table 1 The ET values of FD and KAT parameters in different mole fraction of alcoholic solvent (X2). X2
0.00 0.05 0.10 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60 0.65 0.70 0.75 0.80 0.85 0.90 0.95 1.00
Water–methanol
Water-1-propanol
Water-2-propanol
ET
π*
β
α
ET
Water-ethanol π*
β
α
ET
ET
π*
β
α
58.255 58.184 58.113 58.042 57.983 57.924 57.872 57.825 57.784 57.749 57.720 57.685 57.662 57.638 57.621 57.609 57.598 57.592 57.586 57.575 57.551
1.09 1.14 1.15 1.13 1.10 1.07 1.04 1.01 0.98 0.94 0.91 0.88 0.85 0.82 0.78 0.75 0.72 0.69 0.66 0.63 0.60
0.47 0.54 0.59 0.63 0.65 0.67 0.68 0.69 0.69 0.69 0.68 0.68 0.67 0.67 0.67 0.67 0.68 0.69 0.70 0.70 0.66
1.17 1.14 1.11 1.07 1.05 1.02 1.00 0.99 0.98 0.98 0.98 0.98 0.99 0.99 1.00 1.01 1.01 1.01 1.01 1.01 0.98
58.255 58.066 57.883 57.755 57.656 57.580 57.528 57.482 57.453 57.418 57.390 57.361 57.332 57.303 57.275 57.240 57.217 57.189 57.166 57.137 57.109
1.09 1.16 1.17 1.15 1.10 1.05 0.99 0.93 0.88 0.83 0.79 0.76 0.73 0.70 0.68 0.66 0.65 0.64 0.62 0.59 0.54
0.47 0.44 0.48 0.54 0.60 0.64 0.67 0.68 0.67 0.66 0.65 0.64 0.64 0.65 0.67 0.68 0.70 0.71 0.71 0.71 0.75
1.17 1.14 1.09 1.04 1.00 0.98 0.96 0.96 0.96 0.96 0.96 0.96 0.96 0.95 0.94 0.93 0.92 0.92 0.91 0.91 0.86
58.255 57.907 57.737 57.638 57.563 57.493 57.465 57.424 57.384 57.349 57.315 57.280 57.240 57.206 57.166 57.126 57.080 57.029 56.972 56.910 56.836
58.255 57.924 57.737 57.615 57.534 57.465 57.413 57.367 57.321 57.280 57.235 57.194 57.149 57.103 57.052 56.995 56.932 56.865 56.791 56.707 56.611
1.09 1.07 1.03 0.97 0.92 0.86 0.82 0.78 0.75 0.73 0.72 0.70 0.69 0.68 0.67 0.66 0.64 0.62 0.59 0.55 0.48
0.47 0.54 0.59 0.63 0.65 0.67 0.68 0.69 0.70 0.71 0.71 0.72 0.74 0.75 0.77 0.79 0.81 0.83 0.84 0.83 0.81
1.17 1.08 1.01 0.96 0.93 0.91 0.91 0.91 0.91 0.91 0.91 0.90 0.88 0.86 0.83 0.81 0.79 0.77 0.76 0.76 0.76
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Table 2 Multiple linear regression analysis of the KAT equation for the ET values of FDa. Solvent
A0
Pure solvents
53.53(0.24) 4.06(0.26)
Water– methanol
Water– ethanol
Water-2propanol Fig. 3. The plots of the ET values of FD in aqueous solutions of methanol, ethanol, 1-propanol and 2-propanol as a function of mole fraction of alcoholic solvent at 298.15 K. (Dashed lines show ideal behavior).
means of the multiple linear regressions analysis in Microsoft EXCEL program [48]. Regression was run on all possible combination of three parameters of α, β or π*; three cases including only one of parameters; three cases including two of parameters, and one case including all three of parameters. Fitting results obtained from multiple linear regressions analysis are shown in Table 2; the number in the bracket shows the standard deviation of each regression coefficient; F and r2 represent the value of F-statistic and squared correlation coefficient respectively. The F-statistic is calculated by F¼
ðTSS‐RSSÞ=ðp‐1Þ RSS=ðN−pÞ
ð3Þ
where N and p are the number of data points and fit parameters respectively. The total sum of squares (TSS) and the residual sum of squares (RSS) are defined as TSS ¼
X
Exp
yi
2 −y
ð4Þ
i
RSS ¼
X Exp Cal 2 yi −yi :
ð5Þ
i
Here, yExp is the ith experimental ET value, yCal is the ith ET value cali i culated from a fitted model, and y is the average of the experimental data. A statistical analysis was used to identify which of the single, dualparametric or multi-parametric KAT equations is the best fitted model [49–52]. The regression analysis indicates that the KAT equation with the highest value of F-statistic involves those solvatochromic parameters which have meaning in describing the solvent effect on the obtained results; thus, comparing the models was done using F-statistic values; the larger the F, the more appropriate the model. Accordingly the best fitted-model for each solvent is bolded in Table 2.
59.98(0.42) 55.73 (0.39) 55.24(0.93) 53.31(0.50) 58.91(1.67) 55.11(0.79) 54.36(0.49) 59.82(0.33) 56.81(0.10) 55.55(2.19) 58.20(0.16) 55.16(0.13) 55.98(0.50) 53.61(0.19) 59.66(0.18) 56.35(0.12) 51.61(1.36) 58.39(0.38) 54.20(0.18) 52.76(0.80) 53.90(0.15)
a
b
r2b
p
Fc
0.99 239.87 −3.74(0.56)
0.94 44.17 2.38(0.58) 0.85 17.09 3.02(0.59) −1.04(0.55) 1.00 223.36 4.50(0.89) −0.30(0.56) 0.99 91.19 −2.85(1.48) 0.66(0.99) 0.95 18.18 3.54(0.63) −1.14(0.47) −0.42(0.31) 1.00 209.06 3.35(0.48) 0.72 48.65 −3.10(0.51) 2.65(1.35) −0.72(1.30) −1.71(0.19) 1.91(0.14) 1.43(0.31) −0.50(0.30) 3.96(0.19) −3.44(0.28) 5.23(0.88) 2.99(0.25) 4.52(0.53) 3.77(0.17)
1.18(0.80) −2.23(0.40) 1.46(0.47)
60.17(0.11) −4.09(0.16) 55.49(0.09) 58.36(1.68) 1.09(1.01) −2.92(1.10) 58.37(0.48) −2.55(0.42) 54.55(0.12) 2.10(0.24) 54.19(1.24) 2.27(0.64) 0.25(0.86)
0.66 1.09(0.11) 0.83 0.72 0.79(0.06) 0.97 0.75(0.04) 0.98 0.75(0.04) 0.99 0.96
37.61 92.27 23.60 290.26 558.30 411.70 414.79
0.88 1.33(0.14) 0.82 0.96 0.58(0.16) 0.93 0.42(0.09) 0.98 0.44(0.07) 0.99 0.96
146.07 86.31 221.37 126.47 442.60 439.15 493.68
0.97 2.33(0.12) 0.95 0.97 0.92(0.24) 0.98 1.11(0.15) 0.99 1.15(0.20) 0.99
652.65 383.02 329.73 564.80 978.82 619.28
In pure protic solvents, the fit to the single term α leads to the highest F-statistic value; meaning the spectral changes of FD are mainly attributed to the hydrogen-bond donor acidity of the solvent. The sign of α coefficient is positive; indicating that the electronic transition energy of FD decreases when the hydrogen bond donor ability of solvent decreases. It is worthwhile to note that the sign of π* coefficient is negative in all models including the α term, consistent with the compensation of the negative polarity contribution to the spectral shift of FD by a strong positive contribution of hydrogen bonding interaction of FD with the solvent. In all aqueous binary mixtures, a dual-parametric KAT equation containing α and π* shows the best fit with the lowest standard deviation and the highest F values; thus the solvent effect on the spectral shift of FD can be adequately explained by a combination of α and π* parameters; however the most significant contribution is attributed to α parameters. As expected, the sign of α coefficient is positive in all cases. Interestingly the π* coefficient has also a positive sign in all fitted model. This unexpected result origin is from the different forms of interactions in mixed solvents. Interactions in mixed solvents are more complex than that in pure solvents; they involve not only solute–solvent interactions, but also solvent–solvent interactions. Furthermore the solute molecules can interact to a lesser or greater degree with each of the solvents, giving rise to preferential solvation. 3.2. Preferential solvation of FD in aqueous binary solvents As shown in Fig. 3, the ET value of FD varies in a nonlinear fashion with the mole fraction of organic solvent in all the four binary mixtures. Dashed lines represent the ideal solvation behavior, in which the composition of solvents in the cybotactic region of the solute (local composition) is the same as that in the bulk. A deviation from linearity is clearly evident, resulting in a preferential solvation of FD by one of the components in the mixtures. As observed, in all binary mixtures, the
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deviation from ideality is negative, implying the FD molecules are preferentially solvated by the alcoholic components whose corresponding ET is lower as compared with the water; The extent of preferential solvation decreases in alcoholic-rich region and the ET value reaches approximately to an ideal solvation behavior as the cosolvent changes from methanol to 2-propanol. To explain in greater detail, the variation of the ET value of FD was treated to calculate the extent of preferential solvation by aiming a modified solvent exchange model introduced by El-Seoud [6,19–21]. Considering this model, for a binary mixture containing two solvents mixed S1, S2 and solvatochromic indicator I, the following exchange equilibria can be written for the competition of solvents to solvate the solute. IðS2Þm þ mS1⇌IðS1Þm þ mS2
ð6Þ
IðS2Þm þ mS12⇌IðS12Þm þ mS2
ð7Þ
where I(S1), I(S2) and I(S12) represent the solute solvated by S1, S2 and S12, respectively. The parameter m shows the number of solvent molecules in which their exchange in the solvation shell affects the solvatochromic properties of the solute. Solvating species S12 stands for the hydrogen bonded S1–S2 complexes which are formed by solvent–solvent interactions. The most striking feature of El-Seoud model is that the formation of S12 explicitly is considered in bulk mixtures according to the following equation as the ratio of 1:1. S1 þ S2⇌S12
ð8Þ
The stoichiometric 1:1 model has been verified in aqueous alcoholic mixtures by theoretical and experimental studies [53,54]. The constants of the exchange equilibria in Eqs. (6) and (7) are defined by the preferential solvation parameters (f1/2, f12/2 and f12/1) which relate the mole fraction of solvating species in the solvation shell to that of the same species in bulk. f 1=2
xL ¼ 1L x2
eff
x2
!m ð9Þ
xeff 1 !m
f 12=2
eff xL x2 ¼ 12 L x2 xeff 12
f 12=1
f 12=2 xL12 xeff 1 ¼ ¼ L f 1=2 x1 xeff 12
ð10Þ !m ð11Þ
In fact, fi/j quantifies the tendency of the solute to be solvated by solvent i in the preference of solvent j; the fi/j N 1 means that the solute is more selectively solvated by solvent i relative to solvent j, and vice versa. The xLi and xeff i are the local and effective mole fractions of solvating species of i in the solvation shell and bulk mixture respectively. Effective mole fractions are related, through the association constant (Kassoc) of Eq. (8), to the analytical mole fraction of pure solvents
(X1 and X2). It is clear that the sum of all mole fractions in each region must be equal to unity. eff
eff
eff
L
L
L
X1 þ X2 ¼ x1 þ x2 þ x12 ¼ x1 þ x2 þ x12
ð12Þ
The observed ET value depends directly on the local composition and, as given in Eq. (13), is a weighted average of ET values corresponding to each of solvating species present in the solvation shell. L
L
L
ET ¼ E1 x1 þ E2 x2 þ E12 x12
ð13Þ
Where E1, E2 and E12 are ET values if solute is solvated by pure S1, S2 and S12 respectively. Introducing Eqs. (9), (10) and (12) into Eq. (13) and rearranging in terms of effective mole fractions gives Eq. (11). m m m E2 xeff þ f 1=2 E1 xeff þ f 12=2 E12 xeff 2 1 12 m m m ET ¼ eff eff xeff þ f x þ f x 1=2 12=2 2 1 12
ð14Þ
In this work, S1 and S2 represent pure water and alcoholic solvents respectively. The value of Kassoc is necessary for the calculation of effective mole fractions in each binary mixture. The Kassoc can be calculated from nonideal dependency of some physical properties of binary mixtures, such as density, on their bulk composition [20,21]. The calculated association constant for the formation of S12 in binary mixture of water with methanol, ethanol, 1-propanol and 2-propanol is 0.220, 0.032, 0.012 and 0.015 M−1 respectively [21]. The effective mole fractions were calculated as a function of the analytical mole fraction of alcoholic solvents by considering the density values of each binary mixture extracted from the literature [55,56]. The experimental ET values were fitted to Eq. (14) by aiming a nonlinear regression procedure performed in OriginPro 8.5; results are shown in Table 3. The quality of fits was judged to be excellent, according to low residual sum of squares (RSS) and very good squared correlation coefficient (r2); Also the calculated ET values for pure solvents are in satisfactory agreement with the experimental values in Table 1. Solvating species with higher α acidity tends to interact more strongly with FD, because FD has several hydrogen bond accepting sites. However the preferential solvation parameter f1/2 is lower than unity in all binary mixtures, indicating that the FD is selectively surrounded by the alcohol instead of the water for which the value of α acidity is the maximum one in the mixtures (Table 1). Since the extent of f1/2 increases with the lengthening of the alkyl side of alcohols, it is implausible to attribute the direction of f1/2 to the hydrophobic interactions between FD and alcohols. In fact, the f1/2 b 1 is traced to the unique physicochemical behavior of water; water molecules like to bond to each other by the strong and extensive hydrogen bonding interactions to form a spanning hydrogen-bonded network structure which can not be easily disrupted; it has been also established that the network structure of water molecules is enhanced in the presence of cosolvents such as alcohols [57,58]. Therefore in a competition with the associated structure of water molecules, the alcohol molecules can freely solvate
Table 3 Analysis of preferential solvation model for FD in mixture of water (S1) with methanol, ethanol, 1-propanol and 2-propanol.a S2
E1
E2
E12
m
f1/2
f12/2
f12/1
r2b
RSSc
Methanol Ethanol 1-Propanol 2-Propanol
58.256(0.005) 58.256(0.004) 58.255(0.004) 58.255(0.002)
57.562(0.004) 57.116(0.003) 56.838(0.004) 56.613(0.002)
57.694(0.056) 57.322(0.016) 57.250(0.020) 57.179(0.008)
1.03 1.32 0.99 1.06
0.34 0.50 0.61 0.73
0.99 11.28 14.13 15.45
2.91 22.56 23.16 21.16
0.999 0.999 0.999 0.999
3.56E−04 2.55E−04 2.35E−04 5.62E−05
a b c
Number in bracket is standard deviation. Squared correlation coefficient. Residual of sum squares values.
F. Naderi et al. / Journal of Molecular Liquids 190 (2014) 126–132
the solute by the hydrogen bonding interaction, leading to the f1/2 b 1. Consequently, the cosolvents showing more acidity would be more selectively incorporated into the surrounding shell of solute in preference to the water. Table 3 indicates that the magnitude of f1/2 is in complete agreement with this interpretation. The value of f1/2 increases toward unity in the order methanol b ethanol b 1-propanol b 2-propanol, meaning that the extent of preferential solvation of FD by the alcohol relative to the water diminishes in the binary mixtures, in accordance with the decreasing order of the α acidity of alcohols. This conclusion can adequately explain the trend of the ET value of FD in alcohol-rich region (Fig. 3), which tends to reach an ideal behavior when cosolvent changes from methanol to 2-propanol. The preferential solvation parameter f12/2 is higher than unity in all binary mixtures with the exception of methanol, indicating that the FD is more efficiently solvated by S12 as compared to S2. In this direction, the α acidity of S12 was estimated according to its calculated ET value. As discussed by the KAT analysis, the ET value of FD is significantly sensitive to the hydrogen bond donor acidity of the solvent; the value of ET increases when the hydrogen bond acidity of media increases. Table 3 shows that the ET values of S12 (E12) are higher than the E2 and lower than the E1 in all binary mixtures, exhibiting the α acidity of solvating species increases in the order of alcoholic solvent (S2), solvent complex (S12) and water (S1). Thus, in competition between S2 and S12, the FD molecules more favorably interact with S12 due to a stronger hydrogen bonding interaction, resulting in the f12/2 N 1. With regard to the order of f12/2, the calculated ET values give also useful information. The difference between the E12 and E2 is 0.132, 0.206, 0.412 and 0.566 for the mixture of water with methanol, ethanol, 1-propanol and 2-propanol respectively. This means that, in each binary mixture, the difference between the acidity of S2 and S12 increases as methanol b ethanol b 1propanol b 2-propanol, which corroborates the increasing order of the f12/2 shown in Table 3. In all binary mixtures, f12/1 is higher than unity, indicating that the hydrogen-bonded solvent complex is more efficient in solvating FD than water. In addition, the magnitude of f12/1 is greater than f12/2 in all cases. It is reasonable because as discussed above the water molecules favor the spanning hydrogen-bonded structure, while the methanol molecules make a very fragile network and hence are more willing to compete with S12 in solvating of the solute. In general, the direction of the inequality in f1/2, f12/2 and f12/1 indicates that in water–methanol mixture, the tendency to solvate FD is as S2 ≈ S12 N S1, whereas in other binary mixtures, it is as S12 N S2 N S1. It is interesting to note that all sets of mixtures in Fig. 3, show an inflection point which shifts toward the water-rich regions with the lengthening of the alkyl side of alcohols. This behavior can be explained by the capability of the alcohol to form hydrogen bonded complex (S12) with water. The increase in the alcoholic content disrupts the net structure of water in water-rich region. Since the interaction involving water and alcohol increases with the increase in the β value of alcohol, the formation of S12 and the breaking of net structure take place in the higher mole fraction of water in the order of methanol b ethanol b 2propanol b 1-propanol. 4. Conclusion The solvatochromic properties of FD were successfully studied by UV–vis spectroscopy. Results indicate that the spectral shift of FD is strongly affected by the hydrogen bonding acidity and the polarity of milieu in pure protic and aprotic solvent respectively. In binary mixtures of water with alcohols, both the α and π* parameters are responsible to the spectral changes, however the contribution of the α acidity is most significant. Although, the hydrogen bonding donor acidity of water is highest among the mixture constituents, the downward curvature of the ET value versus the analytical mole fraction of alcohol indicate that the FD is more solvated by the alcoholic components. It was discussed that the self-association of water molecules to make network structure
131
is responsible for the reduction of the water affinity to compete with the other solvating species. The solvent exchange model was perfectly fitted to the observed ET values. The preferential solvation parameters were calculated from the model. The results demonstrate that the solution behavior of FD in the studied mixtures can be adequately explained by the competition of solvating species to interact with the FD through hydrogen bonding interactions. In water–methanol mixture, the affinity of S2 and S12 to the solvation of FD is approximately the same, while in the other mixtures, the solvent complex preferentially solvates the solute. Acknowledgment The authors gratefully acknowledge the financial support from the Research Council of Islamic Azad University Shahr-e Qods branch. References [1] Y. Marcus, Solvent Mixtures: Properties and Selective Salvation, Marcel Dekker, New York, 2002. [2] C. Reichardt, Solvents and Solvent Effects in Organic Chemistry, Wiley-VCH, Weinheim, Germany, 2004. [3] A. Farajtabar, F. Jaberi, F. Gharib, Spectrochim. Acta A 83 (2011) 213. [4] C. Reichardt, Chem. Rev. 94 (1994) 2319. [5] A. Maitra, S. Bagchi, J. Mol. Liq. 137 (2008) 131. [6] O.A. El Seoud, Pure Appl. Chem. 79 (2007) 1135. [7] M. Umadevi, A. Suvitha, K. Latha, B.J.M. Rajkumar, V. Ramakrishnan, Spectrochim. Acta A 67 (2007) 910. [8] L.A. Giusti, V.G. Marini, V.G. Machado, J. Mol. Liq. 150 (2009) 9. [9] A. Ben-Naim, J. Phys. Chem. 93 (1989) 3809. [10] Y. Marcus, J. Chem. Soc. Faraday Trans. 1 (84) (1988) 1465. [11] Y. Marcus, Aust. J. Chem. 36 (1983) 1719. [12] P. Suppan, J. Chem. Soc. Faraday Trans. 1 (83) (1987) 495. [13] O.B. Nagy, M. Wa Muanda, J.B. Nagy, J. Phys. Chem. 83 (1979) 1961. [14] A.K. Covington, K.E. Newman, Pure Appl. Chem. 51 (1979) 2041. [15] U. Buhvestov, F. Rived, C. Ràfols, E. Bosch, M. Rosés, J. Phys. Org. Chem. 11 (1998) 185. [16] S. Hisaindee, J. Graham, M.A. Rauf, M. Nawaz, J. Mol. Liq. 169 (2012) 48. [17] M.L. Moita, R.A. Teodoro, L.M. Pinheiro, J. Mol. Liq. 136 (2007) 15. [18] R.D. Skwierczynski, K.A. Connors, J. Chem. Soc. Perkin Trans. 2 (0) (1994) 467. [19] E.B. Tada, P.L. Silva, C. Tavares, O.A. El Seoud, J. Phys. Org. Chem. 18 (2005) 398. [20] E.B. Tada, P.L. Silva, O.A. El Seoud, J. Phys. Org. Chem. 16 (2003) 691. [21] E.L. Bastos, P.L. Silva, O.A. El Seoud, J. Phys. Chem. A 110 (2006) 10287. [22] S.C. Burdette, G.K. Walkup, B. Spingler, R.Y. Tsien, S.J. Lippard, J. Am. Chem. Soc. 123 (2001) 7831. [23] B.N.G. Giepmans, S.R. Adams, M.H. Ellisman, R.Y. Tsien, Science 312 (2006) 217. [24] F. Hou, J. Cheng, P. Xi, F. Chen, L. Huang, G. Xie, Y. Shi, H. Liu, D. Bai, Z. Zeng, Dalton Trans. 41 (2012) 5799. [25] L.M. Smith, J.Z. Sanders, R.J. Kaiser, P. Hughes, C. Dodd, C.R. Connell, C. Heiner, S.B.H. Kent, L.E. Hood, Nature 321 (1986) 674. [26] H. Wallrabe, A. Periasamy, Curr. Opin. Biotechnol. 16 (2005) 19. [27] V.R. Batistela, J. da Costa Cedran, H.P. Moisés de Oliveira, I.S. Scarminio, L.T. Ueno, A. Eduardo da Hora Machado, N. Hioka, Dyes Pigments 86 (2010) 15. [28] N. Klonis, W. Sawyer, J. Fluoresc. 6 (1996) 147. [29] R. Sjöback, J. Nygren, M. Kubista, Spectrochim. Acta A 51 (1995) L7. [30] R. D'Amato, E. Wesolowski, L.E. Hodgson Smith, Microvasc. Res. 46 (1993) 135. [31] R.W.F. Gross, Jerry F. Bott, Handbook of Chemical Lasers, Wiley, New York, 1976. [32] P.L. Smart, I.M.S. Laidlaw, Water Resour. Res. 13 (1977) 15. [33] M. Ali, P. Dutta, S. Pandey, J. Phys. Chem. B 114 (2010) 15042. [34] S. Biswas, S.C. Bhattacharya, P.K. Sen, S.P. Moulik, J. Photochem. Photobiol. A 123 (1999) 121. [35] M.F. Choi, P. Hawkins, Spectrosc. Lett. 27 (1994) 1049. [36] N. Klonis, A.H.A. Clayton, E.W. Voss, W.H. Sawyer, Photochem. Photobiol. 67 (1998) 500. [37] M.M. Martin, Chem. Phys. Lett. 35 (1975) 105. [38] N.O. McHedlov-Petrossyan, V.V. Ivanov, Russ. J. Phys. Chem. 81 (2007) 112. [39] A. Tamulis, J. Tamuliene, M.L. Balevicius, Z. Rinkevicius, V. Tamulis, Struct. Chem. 14 (2003) 643. [40] P. Zhou, J. Liu, S. Yang, J. Chen, K. Han, G. He, Phys. Chem. Chem. Phys. 14 (2012) 15191. [41] E.A. Slyusareva, F.N. Tomilin, A.G. Sizykh, E.Y. Tankevich, A.A. Kuzubov, S.G. Ovchinnikov, Opt. Spectrosc. 112 (2012) 671. [42] J.M. Alvarez-Pez, L. Ballesteros, E. Talavera, J. Yguerabide, J. Phys. Chem. A 105 (2001) 6320. [43] R. Taft, J.-L. Abboud, M. Kamlet, M. Abraham, J. Solution Chem. 14 (1985) 153. [44] M.J. Kamlet, J.L. Abboud, R.W. Taft, J. Am. Chem. Soc. 99 (1977) 6027. [45] R.W. Taft, M.J. Kamlet, J. Am. Chem. Soc. 98 (1976) 2886. [46] M.J. Kamlet, R.W. Taft, J. Am. Chem. Soc. 98 (1976) 377. [47] Y. Marcus, J. Chem. Soc. Perkin Trans. 2 (1994) 1751. [48] E.J. Billo, Excel for Chemists: a Comprehensive Guide, Wiley, Weinheim, Germany, 2001. [49] F. Naderi, A. Farajtabar, F. Gharib, J. Solution Chem. 41 (2012) 1033.
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Journal of Molecular Liquids journal homepage: www.elsevier.com/locate/molliq
Solvatochromic and preferential solvation of fluorescein in some water-alcoholic mixed solvents Fereshteh Naderi a,⁎, Ali Farajtabar b, Farrokh Gharib c a b c
Department of Chemistry, Shahr-e Qods branch, Islamic Azad University, Tehran, Iran Department of Chemistry, Jouybar branch, Islamic Azad University, Jouybar, Iran Department of Chemistry, Faculty of Sciences, Shahid Beheshti University, Tehran, Evin, Iran
a r t i c l e
i n f o
Article history: Received 27 August 2013 Received in revised form 17 October 2013 Accepted 31 October 2013 Available online 14 November 2013 Keywords: Fluorescein Solvatochromism Preferential solvation Hydrogen bonding
a b s t r a c t It is known that the hydrogen bonding has a significant effect on the spectral properties of fluorescein, however its relevance to solution behavior and preferential solvation of fluorescein is unclear in mixed solvents. In this work, absorption spectra of fluorescein dianion have been investigated in pure water, methanol, ethanol, 1-propanol, 2-propanol, dimethylsulfoxide, N,N-dimethylformamide, acetonitrile and in binary mixtures of water with the mentioned alcohols at 298.15 K. Spectral changes were interpreted in terms of specific and nonspecific solute-solvent interactions. The energy of electronic transition in maximum absorption (ET) was calculated in each binary mixture. The extent and importance of each of solute–solvent interactions to ET were analyzed in the framework of the linear solvation energy relationships. Preferential solvation was detected as a nonideal behavior of ET curve respective to the analytical mole fraction of alcohols in all binary mixtures. The ET values were fitted to a solvent exchange model to calculate the preferential solvation parameters. The preference of fluorescein dianion to be solvated by one of the solvating species relative to others was discussed in terms of solvent–solvent and solute–solvent interactions. © 2013 Elsevier B.V. All rights reserved.
1. Introduction Study on the solvation process opens the way to further understand the solvent effects on the kinetics and thermodynamics of reactions. The behavior of solute in solution is a function of all types of interactions taken place within the solvation shell. Evidently, solvation in mixed solvent is more complicated than pure solvents because of the interference of the solvent–solvent interactions and diversity of solute–solvent interactions. Interaction between solvent molecules can create new solvating complexes whose properties and their structures are distinct and different from those in the pure solvent. Furthermore, the molecules of solute can usually interact to different degrees with each solvent component. At this case, the composition of solvation shell near close to solute, sometimes called local composition, can differ from that of the bulk composition of mixed solvent. Therefore a specific feature of solvation in the mixture of solvents emerges as a result of the phenomenon which is termed as preferential solvation [1,2]. In other words, preferential solvation arises when the solute interacts with more molecules of one solvent component in its cybotactic region than the others in comparison of the bulk composition. According to the above explanation, the comprehension of preferential solvation is a key issue in solvation processes in mixed solvents.
⁎ Corresponding author. Tel./fax: +98 21 46842938. E-mail address: [email protected] (F. Naderi). 0167-7322/$ – see front matter © 2013 Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.molliq.2013.10.028
Solvatochromism, which describes spectral changes of solute induced by changes in solvent polarity, provides a convenient way for the monitoring of interactions that occur in the cybotactic region of solute [2–5]. In particular, the electronic transition energy of the solute in the absorption maxima named as polarity scale, ET, comprises all sorts of interactions between solute and solvent components including specific (such as hydrogen bonding and electron donor acceptor interactions) and nonspecific (polarity–polarizability effects) interactions [2,4,6–8]. The shifts of the absorption maxima with changing the solvent composition and ET values can therefore reveal the nature of the solute–solvent as well as solvent–solvent interactions present in the mixed solvents. Accordingly, ET parameters of a solute can be manipulated in order to elucidate the preferential solvation phenomena in the mixtures. Various theoretical and experimental models to the problem of preferential solvation have been developed in the literature. Between theoretical approaches, the Kirkwood–Buff theory, the quasi lattice-quasi chemical theory developed by Marcus, the dielectric enrichment, the competitive preferential solvation theory of Nagy, and the stepwise solvent exchange model of Covington has been the basis of many papers [9–14]. The solvent exchange model has been proposed by Bosch and Roses to investigate preferential solvation experimentally [15–17]. This model is an extension of the stepwise solvent exchange model of Connors which derives equations that relate ET values of a solute with the solvent composition [18]. In this model, competition between different solvent species to solvate the solute is described by some exchange
F. Naderi et al. / Journal of Molecular Liquids 190 (2014) 126–132
equilibria in vicinity of solute. An equilibrium constant, named preferential solvation parameter, is defined for each exchange reaction, which relates the mole fraction of solvents in solvation shell to that in the bulk mixture. In addition, formation of solvating complexes is postulated from solvent–solvent interaction on the microsphere of solvation. Recently, El-Seoud has modified this model by considering the formation of solvating complexes explicitly in both cybotactic region and bulk mixture, since to subtle description of solvation, all of species in the solvation shell should be considered in equilibrium with the same species in the bulk mixture [19–21]. Fluorescein, a xanthene dye, is a popular fluorescent probe with widespread applications to various areas such as biology, chemistry and technology [22–26]. Fluorescein and fluorescein derivatives have intense absorbance and fluorescence characteristics with high quantum yield in the visible region, low toxicity and the ability of attaching to some biomolecular and ionic species. The presence of seven prototropic forms with different photophysical properties has been identified for fluorescein depending on the pH of the solution [27–29]. Predominant form of fluorescein under neutral and alkaline solutions is fluorescein dianion (FD) which has high fluorescence quantum yield and large absorption coefficient [29]. FD, Fig. 1, is commercially available in sodium salt, known as uranine or resorcinolphthalein sodium. The spectral characteristics of FD are important factors which made it useful in water tracing, laser tuning, ophthalmic photography and fluorescein angiography [30–32]. Accordingly solvatochromic properties of fluorescein have been subject of some studies in aqueous and nonaqueous solvents [33–38]. As reported in literature, position and intensity of absorbance and fluorescence emission maxima in FD are generally dependent on the hydrogen bonding propensity of the solvent. To achieve a further insight into the solution behavior of FD, in this work solvatochromic properties as well as preferential solvation of FD were studied in some pure solvents and binary mixtures at 298.15 K. 2. Experimental 2.1. Materials and methods All pure solvents including methanol, ethanol, 1-propanol, 2propanol, dimethylsulfoxide (DMSO), N,N-dimethylformamide (DMF) and acetonitrile (ACN) were supplied as analytical reagent grade from Merck. Fluorescein disodium and potassium hydroxide were obtained from Sigma. NaOH solution was prepared from a titrisol solution (Merck). All dilute solutions were prepared from double-distilled water with a conductance equal to (1.3 ± 0.1) μS. Spectrophotometric measurements were performed on a UV–vis Shimadzu 2100 spectrophotometer with a Pentium 4 computer and using thermostated matched 10 mm quartz cells at 298.15 K. All aqueous binary mixtures were carefully prepared by weighing at the required molar ratio with an electrical balance accurate to ±0.1 mg. Alkaline solution of solvents was prepared by solving suitable amount of KOH into 50 mL of pure and binary mixed solvents. Stock solution of FD was prepared in ethanol. 50 μL of the stock solution was transferred to -
O
O
127
5 mL glass volumetric tubes. After the evaporation of the ethanol under reduced pressure, 3.5 mL of the pure or the binary solvent was pipetted into the tubes and the mixtures were sonicated to a clear and homogenous solution and then kept in the dark. The final concentration of the solute in the tubes was obtained 6.14 μM. To exclude oxygen from the system, a stream of purified nitrogen gas was passed through solutions prior to each spectrophotometric measurement. The electronic absorption spectra of FD were recorded over the wavelength range of (400–600) nm at a rate of 140 nm min−1 with a slit width of 2 nm. At least three replicate spectrophotometric measurements were done for each solution with an accuracy of ±0.05 nm. 3. Results and discussion 3.1. Solvatochromism of FD in pure and aqueous binary solvents Absorption spectra of FD in pure water, methanol, ethanol, 1propanol and 2-propanol are shown in Fig. 2. In water solution, spectrum of FD exhibits a maximum band at 490.8 nm accompanied with a shoulder located around 470 nm, which assigned to the electronic transition from the ground state (S0) to the first (S1) and second excited singlet state (S2) respectively. The predominant electronic transition is mainly described by the HOMO → LUMO excitation, with a small contribution from HOMO-5 → LUMO [27]. The structures of ground and excited states of FD have been provided in the literature [27,39,40]. These results have indicated that both HOMO and LUMO are localized on the xanthene moiety, and the first excited state (S1), with largest oscillator strength, corresponds mainly to the electron transition from HOMO → LUMO. On going from water to 2-propanol, the position of absorption maximum shifts toward longer wavelengths, with an increase in the absorbance intensity. The spectrum of FD in aqueous solutions of methanol, ethanol, 1-propanol and 2-propanol was also investigated over the entire range of the mole fraction of alcoholic solvent. A regular bathochromic shift was observed as the mole fraction of alcohol increases in the aqueous solution for all binary aqueous mixtures in this work. Quantum chemical calculations have provided evidence that the dipole moment of FD increases upon electronic transition from the ground to excited state in gas phase; the excited state is more polar than the ground state [41]. The polarity effect rules that the more polar excited state is a more stable state in media of high polarity; the reverse is true for the less polar ground state. Therefore, in an opposite direction of what was observed for FD in this work, one could reasonably expect that a hypsochromic shift should emerge as the polarity of media decreases upon a solvent change from water to alcoholic media. This suggests that the nonspecific interactions have no pronounced
O
COO-
Fig. 1. The molecular structure of fluorescein dianion.
Fig. 2. Absorption spectra of FD (6.14 μM) in water, methanol, ethanol, 1-propanol and 2-propanol at 298.15 K.
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media increases due to approximately pure electrostatic interactions. This contradictory is attributed to the difference over the nature of the ground and excited state; the ground state has appreciably a more hydrogen band accepting capacity, while the excited state is a more polar state. It seems that the hydrogen bonding interaction is also responsible for the reduction in the absorption intensity on going from alcoholic media to water, because hydrogen bonding in xanthene ring reduces the resonance of the aromatics cycles. To quantitative representation, the wavelength of the maximum absorbance, λmax, in a particular solvent mixture was calculated from the first derivative of the spectrum data and then transformed into Eq. (1) to calculate the molar electronic transition energies, ET values.
effect on the spectral shift of FD. On the other hand, the FD has several hydrogen bond acceptor sites on its molecular structure, which provide opportunities for the specific solute–solvent interactions. Literature, in consistent with our observation, shows that the electronic spectra of FD experience a blue shift when the hydrogen bond donating power of milieu increases [33–38]. This can be explained by this fact that the acidity of fluorescein in the excited state is significantly greater than that in the ground state, implying that the basicity of FD decreases upon excitation [28,29,42]. Moreover upon photoexcitation, electron densities at the oxygen atoms on the xanthene ring (tricycle) decrease in the excited state. Accordingly, the hydrogen bond accepting nature of FD in excited state is considerably lower than that corresponding to ground sate; thus, as expected, the absorption spectra shift to the shorter wavelength upon hydrogen bonding formation, due to more stabilization of ground state related to the excited state when the hydrogen bond donating power of media increases. The above controversy reveals that the contribution of nonspecific electrostatic solute–solvent interaction to the spectral shift of FD decreases dramatically in the presence of the specific solute–solvent interactions; so the effect of the solvent polarity appears to be masked by the strong hydrogen bonding formation between the FD and the solvent in protic solvents. To gain further verification, the spectral shift of FD was examined in pure aprotic solvents including DMSO, DMF and ACN, in which no significant hydrogen bonding exists between the FD and the solvent, due to the lack of any hydrogen bond donor sites on the FD; seemingly the nonspecific electrostatic interactions appear to be predominant in the absence of the hydrogen bonding interactions. Maximum in the absorption was observed at 515, 519 and 521 nm for ACN, DMF and DMSO respectively, where the polarity of solvent increases as the order ACN b DMF b DMSO. In agreement with the above explanation, a bathochromic shift is clearly seen when the polarity of medium increases, because the dipole–dipole interaction makes the more polar excited state favorably more stable than the ground state. Therefore, in aprotic solvents where hydrogen bond donation is excluded, the spectral shift of FD can be explained in terms of a change in the polarity of the medium. Overall, the net results of hydrogen bonding and electrostatic interactions explain the spectral behavior of FD in solvent. In protic solvents, the hydrogen bonding interactions are dominant for the solute; the main absorption band shifts hypsochromically with increasing the hydrogen bond donating power of media, whereas in aprotic solvents, the absorption bond undergoes a bathochromic shift as the polarity of
−1 ¼ hcNA =λmax ðnmÞ ¼ 28591:5=λmax ðnmÞ ET kcal·mol
ð1Þ
where h, c and NA are Planck's constant, the velocity of light and Avogadro's number, respectively. The values of the ET as a function of mole fraction of alcoholic solvent are presented in Table 1 and Fig. 3. Solvent effect was treated within the framework of the linear solvation energy relationships (LSER) established by Kamlet, Abboud and Taft, in which each of specific and nonspecific interactions have a linear contribution to the total solvation energy of solvent dependent phenomena [43]. In nonchloro-substituted aliphatic solvents, LSER takes the form of Eq. (2). ET ¼ A0 þ aα þ bβ þ pπ
ð2Þ
The α, β and π* are the Kamlet and Taft solvatochromic parameters (KAT) which have been developed for scaling the hydrogen-bond donor acidity, hydrogen-bond acceptor basicity and dipolarity/polarizability of solvents respectively [44–46]. The A0, a, b, and p are regression coefficients; A0 is intercept whereas a, b and p quantify the sensitivity of ET values to the acidity, basicity and dipolarity/polarizability of solvent respectively. The number of parameters in Eq. (2) depends physically on the nature of solute and solvent and the significance of the solute– solvent interactions. The KAT parameters for binary mixtures were extracted from the literature and inserted in Table 1 [47]; the value of α, and π* for 1-propanol is 0.84, 0.90 and 0.54 respectively [2]; no literature data was found for aqueous solutions of 1-propanol. The ET values of FD for each aqueous solution were correlated with solvent properties by
Table 1 The ET values of FD and KAT parameters in different mole fraction of alcoholic solvent (X2). X2
0.00 0.05 0.10 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60 0.65 0.70 0.75 0.80 0.85 0.90 0.95 1.00
Water–methanol
Water-1-propanol
Water-2-propanol
ET
π*
β
α
ET
Water-ethanol π*
β
α
ET
ET
π*
β
α
58.255 58.184 58.113 58.042 57.983 57.924 57.872 57.825 57.784 57.749 57.720 57.685 57.662 57.638 57.621 57.609 57.598 57.592 57.586 57.575 57.551
1.09 1.14 1.15 1.13 1.10 1.07 1.04 1.01 0.98 0.94 0.91 0.88 0.85 0.82 0.78 0.75 0.72 0.69 0.66 0.63 0.60
0.47 0.54 0.59 0.63 0.65 0.67 0.68 0.69 0.69 0.69 0.68 0.68 0.67 0.67 0.67 0.67 0.68 0.69 0.70 0.70 0.66
1.17 1.14 1.11 1.07 1.05 1.02 1.00 0.99 0.98 0.98 0.98 0.98 0.99 0.99 1.00 1.01 1.01 1.01 1.01 1.01 0.98
58.255 58.066 57.883 57.755 57.656 57.580 57.528 57.482 57.453 57.418 57.390 57.361 57.332 57.303 57.275 57.240 57.217 57.189 57.166 57.137 57.109
1.09 1.16 1.17 1.15 1.10 1.05 0.99 0.93 0.88 0.83 0.79 0.76 0.73 0.70 0.68 0.66 0.65 0.64 0.62 0.59 0.54
0.47 0.44 0.48 0.54 0.60 0.64 0.67 0.68 0.67 0.66 0.65 0.64 0.64 0.65 0.67 0.68 0.70 0.71 0.71 0.71 0.75
1.17 1.14 1.09 1.04 1.00 0.98 0.96 0.96 0.96 0.96 0.96 0.96 0.96 0.95 0.94 0.93 0.92 0.92 0.91 0.91 0.86
58.255 57.907 57.737 57.638 57.563 57.493 57.465 57.424 57.384 57.349 57.315 57.280 57.240 57.206 57.166 57.126 57.080 57.029 56.972 56.910 56.836
58.255 57.924 57.737 57.615 57.534 57.465 57.413 57.367 57.321 57.280 57.235 57.194 57.149 57.103 57.052 56.995 56.932 56.865 56.791 56.707 56.611
1.09 1.07 1.03 0.97 0.92 0.86 0.82 0.78 0.75 0.73 0.72 0.70 0.69 0.68 0.67 0.66 0.64 0.62 0.59 0.55 0.48
0.47 0.54 0.59 0.63 0.65 0.67 0.68 0.69 0.70 0.71 0.71 0.72 0.74 0.75 0.77 0.79 0.81 0.83 0.84 0.83 0.81
1.17 1.08 1.01 0.96 0.93 0.91 0.91 0.91 0.91 0.91 0.91 0.90 0.88 0.86 0.83 0.81 0.79 0.77 0.76 0.76 0.76
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Table 2 Multiple linear regression analysis of the KAT equation for the ET values of FDa. Solvent
A0
Pure solvents
53.53(0.24) 4.06(0.26)
Water– methanol
Water– ethanol
Water-2propanol Fig. 3. The plots of the ET values of FD in aqueous solutions of methanol, ethanol, 1-propanol and 2-propanol as a function of mole fraction of alcoholic solvent at 298.15 K. (Dashed lines show ideal behavior).
means of the multiple linear regressions analysis in Microsoft EXCEL program [48]. Regression was run on all possible combination of three parameters of α, β or π*; three cases including only one of parameters; three cases including two of parameters, and one case including all three of parameters. Fitting results obtained from multiple linear regressions analysis are shown in Table 2; the number in the bracket shows the standard deviation of each regression coefficient; F and r2 represent the value of F-statistic and squared correlation coefficient respectively. The F-statistic is calculated by F¼
ðTSS‐RSSÞ=ðp‐1Þ RSS=ðN−pÞ
ð3Þ
where N and p are the number of data points and fit parameters respectively. The total sum of squares (TSS) and the residual sum of squares (RSS) are defined as TSS ¼
X
Exp
yi
2 −y
ð4Þ
i
RSS ¼
X Exp Cal 2 yi −yi :
ð5Þ
i
Here, yExp is the ith experimental ET value, yCal is the ith ET value cali i culated from a fitted model, and y is the average of the experimental data. A statistical analysis was used to identify which of the single, dualparametric or multi-parametric KAT equations is the best fitted model [49–52]. The regression analysis indicates that the KAT equation with the highest value of F-statistic involves those solvatochromic parameters which have meaning in describing the solvent effect on the obtained results; thus, comparing the models was done using F-statistic values; the larger the F, the more appropriate the model. Accordingly the best fitted-model for each solvent is bolded in Table 2.
59.98(0.42) 55.73 (0.39) 55.24(0.93) 53.31(0.50) 58.91(1.67) 55.11(0.79) 54.36(0.49) 59.82(0.33) 56.81(0.10) 55.55(2.19) 58.20(0.16) 55.16(0.13) 55.98(0.50) 53.61(0.19) 59.66(0.18) 56.35(0.12) 51.61(1.36) 58.39(0.38) 54.20(0.18) 52.76(0.80) 53.90(0.15)
a
b
r2b
p
Fc
0.99 239.87 −3.74(0.56)
0.94 44.17 2.38(0.58) 0.85 17.09 3.02(0.59) −1.04(0.55) 1.00 223.36 4.50(0.89) −0.30(0.56) 0.99 91.19 −2.85(1.48) 0.66(0.99) 0.95 18.18 3.54(0.63) −1.14(0.47) −0.42(0.31) 1.00 209.06 3.35(0.48) 0.72 48.65 −3.10(0.51) 2.65(1.35) −0.72(1.30) −1.71(0.19) 1.91(0.14) 1.43(0.31) −0.50(0.30) 3.96(0.19) −3.44(0.28) 5.23(0.88) 2.99(0.25) 4.52(0.53) 3.77(0.17)
1.18(0.80) −2.23(0.40) 1.46(0.47)
60.17(0.11) −4.09(0.16) 55.49(0.09) 58.36(1.68) 1.09(1.01) −2.92(1.10) 58.37(0.48) −2.55(0.42) 54.55(0.12) 2.10(0.24) 54.19(1.24) 2.27(0.64) 0.25(0.86)
0.66 1.09(0.11) 0.83 0.72 0.79(0.06) 0.97 0.75(0.04) 0.98 0.75(0.04) 0.99 0.96
37.61 92.27 23.60 290.26 558.30 411.70 414.79
0.88 1.33(0.14) 0.82 0.96 0.58(0.16) 0.93 0.42(0.09) 0.98 0.44(0.07) 0.99 0.96
146.07 86.31 221.37 126.47 442.60 439.15 493.68
0.97 2.33(0.12) 0.95 0.97 0.92(0.24) 0.98 1.11(0.15) 0.99 1.15(0.20) 0.99
652.65 383.02 329.73 564.80 978.82 619.28
In pure protic solvents, the fit to the single term α leads to the highest F-statistic value; meaning the spectral changes of FD are mainly attributed to the hydrogen-bond donor acidity of the solvent. The sign of α coefficient is positive; indicating that the electronic transition energy of FD decreases when the hydrogen bond donor ability of solvent decreases. It is worthwhile to note that the sign of π* coefficient is negative in all models including the α term, consistent with the compensation of the negative polarity contribution to the spectral shift of FD by a strong positive contribution of hydrogen bonding interaction of FD with the solvent. In all aqueous binary mixtures, a dual-parametric KAT equation containing α and π* shows the best fit with the lowest standard deviation and the highest F values; thus the solvent effect on the spectral shift of FD can be adequately explained by a combination of α and π* parameters; however the most significant contribution is attributed to α parameters. As expected, the sign of α coefficient is positive in all cases. Interestingly the π* coefficient has also a positive sign in all fitted model. This unexpected result origin is from the different forms of interactions in mixed solvents. Interactions in mixed solvents are more complex than that in pure solvents; they involve not only solute–solvent interactions, but also solvent–solvent interactions. Furthermore the solute molecules can interact to a lesser or greater degree with each of the solvents, giving rise to preferential solvation. 3.2. Preferential solvation of FD in aqueous binary solvents As shown in Fig. 3, the ET value of FD varies in a nonlinear fashion with the mole fraction of organic solvent in all the four binary mixtures. Dashed lines represent the ideal solvation behavior, in which the composition of solvents in the cybotactic region of the solute (local composition) is the same as that in the bulk. A deviation from linearity is clearly evident, resulting in a preferential solvation of FD by one of the components in the mixtures. As observed, in all binary mixtures, the
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F. Naderi et al. / Journal of Molecular Liquids 190 (2014) 126–132
deviation from ideality is negative, implying the FD molecules are preferentially solvated by the alcoholic components whose corresponding ET is lower as compared with the water; The extent of preferential solvation decreases in alcoholic-rich region and the ET value reaches approximately to an ideal solvation behavior as the cosolvent changes from methanol to 2-propanol. To explain in greater detail, the variation of the ET value of FD was treated to calculate the extent of preferential solvation by aiming a modified solvent exchange model introduced by El-Seoud [6,19–21]. Considering this model, for a binary mixture containing two solvents mixed S1, S2 and solvatochromic indicator I, the following exchange equilibria can be written for the competition of solvents to solvate the solute. IðS2Þm þ mS1⇌IðS1Þm þ mS2
ð6Þ
IðS2Þm þ mS12⇌IðS12Þm þ mS2
ð7Þ
where I(S1), I(S2) and I(S12) represent the solute solvated by S1, S2 and S12, respectively. The parameter m shows the number of solvent molecules in which their exchange in the solvation shell affects the solvatochromic properties of the solute. Solvating species S12 stands for the hydrogen bonded S1–S2 complexes which are formed by solvent–solvent interactions. The most striking feature of El-Seoud model is that the formation of S12 explicitly is considered in bulk mixtures according to the following equation as the ratio of 1:1. S1 þ S2⇌S12
ð8Þ
The stoichiometric 1:1 model has been verified in aqueous alcoholic mixtures by theoretical and experimental studies [53,54]. The constants of the exchange equilibria in Eqs. (6) and (7) are defined by the preferential solvation parameters (f1/2, f12/2 and f12/1) which relate the mole fraction of solvating species in the solvation shell to that of the same species in bulk. f 1=2
xL ¼ 1L x2
eff
x2
!m ð9Þ
xeff 1 !m
f 12=2
eff xL x2 ¼ 12 L x2 xeff 12
f 12=1
f 12=2 xL12 xeff 1 ¼ ¼ L f 1=2 x1 xeff 12
ð10Þ !m ð11Þ
In fact, fi/j quantifies the tendency of the solute to be solvated by solvent i in the preference of solvent j; the fi/j N 1 means that the solute is more selectively solvated by solvent i relative to solvent j, and vice versa. The xLi and xeff i are the local and effective mole fractions of solvating species of i in the solvation shell and bulk mixture respectively. Effective mole fractions are related, through the association constant (Kassoc) of Eq. (8), to the analytical mole fraction of pure solvents
(X1 and X2). It is clear that the sum of all mole fractions in each region must be equal to unity. eff
eff
eff
L
L
L
X1 þ X2 ¼ x1 þ x2 þ x12 ¼ x1 þ x2 þ x12
ð12Þ
The observed ET value depends directly on the local composition and, as given in Eq. (13), is a weighted average of ET values corresponding to each of solvating species present in the solvation shell. L
L
L
ET ¼ E1 x1 þ E2 x2 þ E12 x12
ð13Þ
Where E1, E2 and E12 are ET values if solute is solvated by pure S1, S2 and S12 respectively. Introducing Eqs. (9), (10) and (12) into Eq. (13) and rearranging in terms of effective mole fractions gives Eq. (11). m m m E2 xeff þ f 1=2 E1 xeff þ f 12=2 E12 xeff 2 1 12 m m m ET ¼ eff eff xeff þ f x þ f x 1=2 12=2 2 1 12
ð14Þ
In this work, S1 and S2 represent pure water and alcoholic solvents respectively. The value of Kassoc is necessary for the calculation of effective mole fractions in each binary mixture. The Kassoc can be calculated from nonideal dependency of some physical properties of binary mixtures, such as density, on their bulk composition [20,21]. The calculated association constant for the formation of S12 in binary mixture of water with methanol, ethanol, 1-propanol and 2-propanol is 0.220, 0.032, 0.012 and 0.015 M−1 respectively [21]. The effective mole fractions were calculated as a function of the analytical mole fraction of alcoholic solvents by considering the density values of each binary mixture extracted from the literature [55,56]. The experimental ET values were fitted to Eq. (14) by aiming a nonlinear regression procedure performed in OriginPro 8.5; results are shown in Table 3. The quality of fits was judged to be excellent, according to low residual sum of squares (RSS) and very good squared correlation coefficient (r2); Also the calculated ET values for pure solvents are in satisfactory agreement with the experimental values in Table 1. Solvating species with higher α acidity tends to interact more strongly with FD, because FD has several hydrogen bond accepting sites. However the preferential solvation parameter f1/2 is lower than unity in all binary mixtures, indicating that the FD is selectively surrounded by the alcohol instead of the water for which the value of α acidity is the maximum one in the mixtures (Table 1). Since the extent of f1/2 increases with the lengthening of the alkyl side of alcohols, it is implausible to attribute the direction of f1/2 to the hydrophobic interactions between FD and alcohols. In fact, the f1/2 b 1 is traced to the unique physicochemical behavior of water; water molecules like to bond to each other by the strong and extensive hydrogen bonding interactions to form a spanning hydrogen-bonded network structure which can not be easily disrupted; it has been also established that the network structure of water molecules is enhanced in the presence of cosolvents such as alcohols [57,58]. Therefore in a competition with the associated structure of water molecules, the alcohol molecules can freely solvate
Table 3 Analysis of preferential solvation model for FD in mixture of water (S1) with methanol, ethanol, 1-propanol and 2-propanol.a S2
E1
E2
E12
m
f1/2
f12/2
f12/1
r2b
RSSc
Methanol Ethanol 1-Propanol 2-Propanol
58.256(0.005) 58.256(0.004) 58.255(0.004) 58.255(0.002)
57.562(0.004) 57.116(0.003) 56.838(0.004) 56.613(0.002)
57.694(0.056) 57.322(0.016) 57.250(0.020) 57.179(0.008)
1.03 1.32 0.99 1.06
0.34 0.50 0.61 0.73
0.99 11.28 14.13 15.45
2.91 22.56 23.16 21.16
0.999 0.999 0.999 0.999
3.56E−04 2.55E−04 2.35E−04 5.62E−05
a b c
Number in bracket is standard deviation. Squared correlation coefficient. Residual of sum squares values.
F. Naderi et al. / Journal of Molecular Liquids 190 (2014) 126–132
the solute by the hydrogen bonding interaction, leading to the f1/2 b 1. Consequently, the cosolvents showing more acidity would be more selectively incorporated into the surrounding shell of solute in preference to the water. Table 3 indicates that the magnitude of f1/2 is in complete agreement with this interpretation. The value of f1/2 increases toward unity in the order methanol b ethanol b 1-propanol b 2-propanol, meaning that the extent of preferential solvation of FD by the alcohol relative to the water diminishes in the binary mixtures, in accordance with the decreasing order of the α acidity of alcohols. This conclusion can adequately explain the trend of the ET value of FD in alcohol-rich region (Fig. 3), which tends to reach an ideal behavior when cosolvent changes from methanol to 2-propanol. The preferential solvation parameter f12/2 is higher than unity in all binary mixtures with the exception of methanol, indicating that the FD is more efficiently solvated by S12 as compared to S2. In this direction, the α acidity of S12 was estimated according to its calculated ET value. As discussed by the KAT analysis, the ET value of FD is significantly sensitive to the hydrogen bond donor acidity of the solvent; the value of ET increases when the hydrogen bond acidity of media increases. Table 3 shows that the ET values of S12 (E12) are higher than the E2 and lower than the E1 in all binary mixtures, exhibiting the α acidity of solvating species increases in the order of alcoholic solvent (S2), solvent complex (S12) and water (S1). Thus, in competition between S2 and S12, the FD molecules more favorably interact with S12 due to a stronger hydrogen bonding interaction, resulting in the f12/2 N 1. With regard to the order of f12/2, the calculated ET values give also useful information. The difference between the E12 and E2 is 0.132, 0.206, 0.412 and 0.566 for the mixture of water with methanol, ethanol, 1-propanol and 2-propanol respectively. This means that, in each binary mixture, the difference between the acidity of S2 and S12 increases as methanol b ethanol b 1propanol b 2-propanol, which corroborates the increasing order of the f12/2 shown in Table 3. In all binary mixtures, f12/1 is higher than unity, indicating that the hydrogen-bonded solvent complex is more efficient in solvating FD than water. In addition, the magnitude of f12/1 is greater than f12/2 in all cases. It is reasonable because as discussed above the water molecules favor the spanning hydrogen-bonded structure, while the methanol molecules make a very fragile network and hence are more willing to compete with S12 in solvating of the solute. In general, the direction of the inequality in f1/2, f12/2 and f12/1 indicates that in water–methanol mixture, the tendency to solvate FD is as S2 ≈ S12 N S1, whereas in other binary mixtures, it is as S12 N S2 N S1. It is interesting to note that all sets of mixtures in Fig. 3, show an inflection point which shifts toward the water-rich regions with the lengthening of the alkyl side of alcohols. This behavior can be explained by the capability of the alcohol to form hydrogen bonded complex (S12) with water. The increase in the alcoholic content disrupts the net structure of water in water-rich region. Since the interaction involving water and alcohol increases with the increase in the β value of alcohol, the formation of S12 and the breaking of net structure take place in the higher mole fraction of water in the order of methanol b ethanol b 2propanol b 1-propanol. 4. Conclusion The solvatochromic properties of FD were successfully studied by UV–vis spectroscopy. Results indicate that the spectral shift of FD is strongly affected by the hydrogen bonding acidity and the polarity of milieu in pure protic and aprotic solvent respectively. In binary mixtures of water with alcohols, both the α and π* parameters are responsible to the spectral changes, however the contribution of the α acidity is most significant. Although, the hydrogen bonding donor acidity of water is highest among the mixture constituents, the downward curvature of the ET value versus the analytical mole fraction of alcohol indicate that the FD is more solvated by the alcoholic components. It was discussed that the self-association of water molecules to make network structure
131
is responsible for the reduction of the water affinity to compete with the other solvating species. The solvent exchange model was perfectly fitted to the observed ET values. The preferential solvation parameters were calculated from the model. The results demonstrate that the solution behavior of FD in the studied mixtures can be adequately explained by the competition of solvating species to interact with the FD through hydrogen bonding interactions. In water–methanol mixture, the affinity of S2 and S12 to the solvation of FD is approximately the same, while in the other mixtures, the solvent complex preferentially solvates the solute. Acknowledgment The authors gratefully acknowledge the financial support from the Research Council of Islamic Azad University Shahr-e Qods branch. References [1] Y. Marcus, Solvent Mixtures: Properties and Selective Salvation, Marcel Dekker, New York, 2002. [2] C. Reichardt, Solvents and Solvent Effects in Organic Chemistry, Wiley-VCH, Weinheim, Germany, 2004. [3] A. Farajtabar, F. Jaberi, F. Gharib, Spectrochim. Acta A 83 (2011) 213. [4] C. Reichardt, Chem. Rev. 94 (1994) 2319. [5] A. Maitra, S. Bagchi, J. Mol. Liq. 137 (2008) 131. [6] O.A. El Seoud, Pure Appl. Chem. 79 (2007) 1135. [7] M. Umadevi, A. Suvitha, K. Latha, B.J.M. Rajkumar, V. Ramakrishnan, Spectrochim. Acta A 67 (2007) 910. [8] L.A. Giusti, V.G. Marini, V.G. Machado, J. Mol. Liq. 150 (2009) 9. [9] A. Ben-Naim, J. Phys. Chem. 93 (1989) 3809. [10] Y. Marcus, J. Chem. Soc. Faraday Trans. 1 (84) (1988) 1465. [11] Y. Marcus, Aust. J. Chem. 36 (1983) 1719. [12] P. Suppan, J. Chem. Soc. Faraday Trans. 1 (83) (1987) 495. [13] O.B. Nagy, M. Wa Muanda, J.B. Nagy, J. Phys. Chem. 83 (1979) 1961. [14] A.K. Covington, K.E. Newman, Pure Appl. Chem. 51 (1979) 2041. [15] U. Buhvestov, F. Rived, C. Ràfols, E. Bosch, M. Rosés, J. Phys. Org. Chem. 11 (1998) 185. [16] S. Hisaindee, J. Graham, M.A. Rauf, M. Nawaz, J. Mol. Liq. 169 (2012) 48. [17] M.L. Moita, R.A. Teodoro, L.M. Pinheiro, J. Mol. Liq. 136 (2007) 15. [18] R.D. Skwierczynski, K.A. Connors, J. Chem. Soc. Perkin Trans. 2 (0) (1994) 467. [19] E.B. Tada, P.L. Silva, C. Tavares, O.A. El Seoud, J. Phys. Org. Chem. 18 (2005) 398. [20] E.B. Tada, P.L. Silva, O.A. El Seoud, J. Phys. Org. Chem. 16 (2003) 691. [21] E.L. Bastos, P.L. Silva, O.A. El Seoud, J. Phys. Chem. A 110 (2006) 10287. [22] S.C. Burdette, G.K. Walkup, B. Spingler, R.Y. Tsien, S.J. Lippard, J. Am. Chem. Soc. 123 (2001) 7831. [23] B.N.G. Giepmans, S.R. Adams, M.H. Ellisman, R.Y. Tsien, Science 312 (2006) 217. [24] F. Hou, J. Cheng, P. Xi, F. Chen, L. Huang, G. Xie, Y. Shi, H. Liu, D. Bai, Z. Zeng, Dalton Trans. 41 (2012) 5799. [25] L.M. Smith, J.Z. Sanders, R.J. Kaiser, P. Hughes, C. Dodd, C.R. Connell, C. Heiner, S.B.H. Kent, L.E. Hood, Nature 321 (1986) 674. [26] H. Wallrabe, A. Periasamy, Curr. Opin. Biotechnol. 16 (2005) 19. [27] V.R. Batistela, J. da Costa Cedran, H.P. Moisés de Oliveira, I.S. Scarminio, L.T. Ueno, A. Eduardo da Hora Machado, N. Hioka, Dyes Pigments 86 (2010) 15. [28] N. Klonis, W. Sawyer, J. Fluoresc. 6 (1996) 147. [29] R. Sjöback, J. Nygren, M. Kubista, Spectrochim. Acta A 51 (1995) L7. [30] R. D'Amato, E. Wesolowski, L.E. Hodgson Smith, Microvasc. Res. 46 (1993) 135. [31] R.W.F. Gross, Jerry F. Bott, Handbook of Chemical Lasers, Wiley, New York, 1976. [32] P.L. Smart, I.M.S. Laidlaw, Water Resour. Res. 13 (1977) 15. [33] M. Ali, P. Dutta, S. Pandey, J. Phys. Chem. B 114 (2010) 15042. [34] S. Biswas, S.C. Bhattacharya, P.K. Sen, S.P. Moulik, J. Photochem. Photobiol. A 123 (1999) 121. [35] M.F. Choi, P. Hawkins, Spectrosc. Lett. 27 (1994) 1049. [36] N. Klonis, A.H.A. Clayton, E.W. Voss, W.H. Sawyer, Photochem. Photobiol. 67 (1998) 500. [37] M.M. Martin, Chem. Phys. Lett. 35 (1975) 105. [38] N.O. McHedlov-Petrossyan, V.V. Ivanov, Russ. J. Phys. Chem. 81 (2007) 112. [39] A. Tamulis, J. Tamuliene, M.L. Balevicius, Z. Rinkevicius, V. Tamulis, Struct. Chem. 14 (2003) 643. [40] P. Zhou, J. Liu, S. Yang, J. Chen, K. Han, G. He, Phys. Chem. Chem. Phys. 14 (2012) 15191. [41] E.A. Slyusareva, F.N. Tomilin, A.G. Sizykh, E.Y. Tankevich, A.A. Kuzubov, S.G. Ovchinnikov, Opt. Spectrosc. 112 (2012) 671. [42] J.M. Alvarez-Pez, L. Ballesteros, E. Talavera, J. Yguerabide, J. Phys. Chem. A 105 (2001) 6320. [43] R. Taft, J.-L. Abboud, M. Kamlet, M. Abraham, J. Solution Chem. 14 (1985) 153. [44] M.J. Kamlet, J.L. Abboud, R.W. Taft, J. Am. Chem. Soc. 99 (1977) 6027. [45] R.W. Taft, M.J. Kamlet, J. Am. Chem. Soc. 98 (1976) 2886. [46] M.J. Kamlet, R.W. Taft, J. Am. Chem. Soc. 98 (1976) 377. [47] Y. Marcus, J. Chem. Soc. Perkin Trans. 2 (1994) 1751. [48] E.J. Billo, Excel for Chemists: a Comprehensive Guide, Wiley, Weinheim, Germany, 2001. [49] F. Naderi, A. Farajtabar, F. Gharib, J. Solution Chem. 41 (2012) 1033.
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