Fluorescence of 8-Quinolinol in Strongly Basic Solutions A Defense of the Excited State Prototropic Equilibrium Hypothesis SIR: 8-Quinolinol and several of its derivatives have been reported to be fluorescent in concentrated acid media but not in dilute acid, neutral, or basic aqueous solutions (1-3). The quenching of the green fluorescence of the 8-quinolinolium ion in solutions of low acidity has been attributed to a prototropic equilibrium in the lowest excited singlet state between the 8-quinolinolium cation and the nonfluorescent 8-quinolinol zwitterion. More recently, the excited state equilibrium hypothesis has been disputed by Goldman and Wehry ( 4 ) who have studied the fluorescence of 8-quinolinol in several media and who concluded that the quenching of 8quinolinol fluorescence is due to hydrogen bonding by hydroxylic solvents. Recently, in this laboratory, green fluorescence has been observed from the 8-quinolinolate anion in carbonate free, nitrogen purged, concentrated sodium hydroxide solutions in water, Fluorescence spectra were taken on a Farrand Mark I Spectrofluorometer whose monochromators were calibrated against the mercury lines from a Pen-Ray low pressure mercury lamp. Emission spectra were not corrected for the wavelength dependence of the monochromators or the phototube, but this correction is not vital to this work. Fluorimetric titration of the 8-quinolinolate anion luminescence, by plotting the relative fluorescence intensity (Z/Zo)at the emission maximum, 483 nm, us. H-, the Hammett acidity function for strongly basic media (5),is represented in Figure 1. This titration curve allows an estimate of the pK* corresponding to equilibrium, in the lowest excited singlet state, between the anion and the zwitterion derived from 8-quinolinol. The latter dissociation constant is estimated to be 15.6. Ballard and Edwards ( I ) have calculated the approximate pK* values for all prototropic reactions of 8-quinolinol by the Forster cycle method (6), using the long wavelength absorption maxima of the cation, neutral, zwitterion, and anion forms of 8-quinolinol along with the appropriate ground state equilibrium constants. These pK* values are not quantitatively accurate since they do not take into account differences in vibrational structures of the ground and excited states of the various prototropic forms of 8-quinolinol, or the differences in solvent cage between the ground and excited state species. In effect, they represent that part of the pK* which is due to the electronic dipole changes produced by the excitation process. However, since 8-quinolinol does not fluoresce in solutions of low acidity or basicity, the pK* values derived from absorption spectra alone, are the only means available, in this case, of comparing Forster cycle data with fluorimetric titration data. Furthermore, if the differences in pK values between the ground and excited states, as determined from absorption data alone, are great (as they are in this case), they will fairly accurately reflect whether the molecule in question becomes more acidic or more basic upon excitation (Le., the effect of the shift in electronic charge, upon excitation, will be greater than the effect of solvent R. E. Ballard and J. Edwards, J . C/iem. Soc., 1964,4868. S. Schulman and Q. Fernando, Tetraliedroil, 24, 1777 (1968). S. Schulman and Q. Fernando, J . Pliys. Clieni., 71, 2668 (1967). M. Goldman and E. L. Wehry, ANAL.CHEM., 42, 1178 (1970). ( 5 ) C. M. Rochester, Qimrt. Reu., Chcm. Soc., 20, 511 (1966). (6) T. Forster, Z . Elektroc/ieni., 54,42 (1950).
(1) (2) (3) (4)
H,
Figure 1. Fluorimetric titration of 1.00 X lO-4M 8-hydroxyquinoline in concentrated NaOH solutions I / I , = Relative fluorescence intensity, H- = Hammett acidity function for basic solutions
relaxation upon the true pK* value of the excited molecule). The values of pK* calculated by Ballard and Edwards, for 8quinolinol prototropic equilibria, indicate that the 8-quinolinolium cation becomes much more acidic in the lowest excited singlet state with respect to dissociation of the phenolic proton, and much less acidic with respect to dissociation of the protonated ring nitrogen. In addition, the anion, in the lowest excited singlet state is much more basic with respect to the ring nitrogen and much less basic with respect to the phenolic oxygen, than in the ground state. These results infer than in neutral, weakly acidic, and weakly basic media, the zwitterion will be the predominant uncharged species in solution, in the lowest excited singlet state, and that protonation of the excited zwitterion will occur at much lower pH (or Hammett acidities) while dissociation of the excited zwitterion will occur at much higher pH (or Hammett acidities) than in the ground state. Ballard and Edwards ( I ) , have also studied the fluorescence and fluorimetric titration behavior of N-methyl-8-quinolinol and 8-methoxyquinoline. The latter molecule in acid solution is similar in electronic structure to the 8-quinolinolium ion, while in basic solution it is electronically similar to the neutral 8-quinolinol. N-Methyl-8-quinolino1, on the other hand, is similar to 8-quinolinium ion in acid solution while in basic solution it is similar to the 8-quinolinol zwitterion. 8-Methoxyquinoline was found to exhibit intense green fluorescence in concentrated and dilute acid solutions and a weak blue-green fluorescence in basic media. The Nmethyl compound however, fluoresced green, only in concen-
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trated acid and did not fluoresce at all in dilute acid, neutral, or basic media. In this laboratory 8-methoxyquinoline was found to exhibit a weak blue-green fluorescence in strongly basic media which was similar in appearance to that of the same compound in dilute basic solution. The similarity in fluorimetric titration behavior of Nmethyl-8-quinolinolium ion with that of 8-quinolino1, the failure of 8-methoxyquinoline to show green fluorescence in concentrated sodium hydroxide solution, the failure of Nmethyl-8-quinolinol to fluoresce in dilute acid, neutral, or basic solutions, and the Forster cycle calculations of Ballard and Edwards, all lend support to prototropic equilibrium in the lowest excited singlet state being responsible for the quenching of the green 8-quinolinium ion fluorescence in concentrated acid solution and the green 8-quinolinolate anion fluorescence in concentrated basic solution. The excited zwitterion appears to be the conjugate base in the former case and the conjugate acid in the latter case. The observation of blue-green fluorescence from 8-methoxyquinoline in basic aqueous solutions is in agreement with the conclusion by Goldman and Wehry (4) that hydrogenbonding, more specifically hydrogen bonding involving the phenolic group as donor, is significant in the quenching of fluorescence from the neutral species, derived from 8-quinolinol, in water. However, the quenching of the fluorescence of the neutral species in ethanol by addition of water cannot be taken as proof positive that hydrogen bonding is actually the mechanism of quenching here. Ethanol is, after all, a hydrogen bonding solvent. In both water and ethanol, the neutral species is the predominant ground state form of 8quinolinol; however, in the lowest excited singlet state, the zwitterion is known t o predominate in water ( I ) . Ethanol, however, has a substantially lower dielectric constant than water and may not stabilize the excited zwitterion. Consequently, the quenching may be due t o tautomerism in the excited state resulting in the formation of the nonfluorescent zwitterion. Comparison of the quenching of the neutral species’ fluorescence in ethanol by addition of water cannot be made with the quenching of the 8-quinolinolium cation fluorescence in concentrated acid by addition of water. In ethanol, whether hydrogen-bonding or tautomerism is the quenching mechanism, no loss or gain of a proton is involved. In acid solution, however, quenching is accompanied by loss of a proton from the excited cation. Accordingly, the fluorimetric titration curve in acid has the appearance of an acid base titration curve while the quenching in ethanol does not. The data of the fluorimetric titration curves fit the Henderson-Hasselbach equation while the quenching data in ethanol d o not. Additionally, the pK* value for 8-quinolinol is the same ( I , 2), whether determined in concentrated sulfuric acid or in concentrated perchloric acid even though these media have different molar proportions of acid/water (7). The concentrated acids, incidentally, are hydroxylic (although acidic) solvents and d o contain substantial amounts of water (as d o the NaOH solutions). It would appear that the protonating ability of the solvent, and not the fraction of water present, per se, is important. The failure of the 8quinolinolium cation to fluoresce at Hammett acidities more positive than, say -4, even though the cation is the predominant species below pH 5.1, and that of the 8-quinolinolate anion to fluoresce at pH below 13 even though the anion predominates above pH 10, are consequences of the thermodynamics of the excited state prototropic reactions and the
favorable rates of proton transfer within the lifetimes of the excited states of the species concerned. Excitation of the cation or anion in a region of acidity in which the zwitterion predominates in the excited state, is followed by dissociation or protonation, respectively, resulting in the formation of the excited zwitterion which rapidly undergoes internal conversion or specific quenching by the solvent and is radiationlessly returned t o the ground state. The fluorescence of the 8-quinolinolium cation in isopentane saturated with HCI and the quenching thereof by the addition of small amounts of ethanol are not in conflict with the excited state prototropic equilibrium hypothesis. In the hydrocarbon solvent, dissociation of the excited cation is impossible so that fluorescence is observed. However, addition of small amounts of the relatively polar and hydrogen bonding ethanol results in “iceberg” formation, with the ethanol solvating the cation molecules. This interaction provides a basic acceptor molecule (the ethanol) t o receive a proton from the highly acidic excited cation resulting in quenching of the green cation fluorescence. The cation still predominating in the ground state, the absorption spectrum should not be expected to change much with the addition of small amounts of ethanol. Since the quantum yield of fluorescence of the cation in isopentane is much greater than that of 8-quinolinol in either n-hexane or in ethanol, no change in shape or position of the emission spectrum is to be expected. It is quite possible, however, that the quenching of the 8-quinolinolium fluorescence in isopentane, by ethanol is explicable, at least in part, on the basis of hydrogen bonding in the excited state between the cation and the ethanol. The quenching of the neutral 8-quinolinol fluorescence, in methanol, by both acid and base, was studied by Goldman and Wehry and yielded titration curves which displayed the acidbase chemistry of the neutral species in the ground state. These static quenching phenomena show the presence of the excited neutral species in methanol and indicate that in near neutral methanol solutions protonation and dissociation of the excited neutral species is too slow to compete with fluorescence for deactivation of the lowest excited singlet state. However, since the dielectric strength of methanol is intermediate between those of water and ethanol, it is possible that a substantial fraction of 8-quinolinol is present in near neutral methanolic solutions, in the excited state, as the zwitterion. This author has observed green fluorescence (presumably from the excited 8-quinolinolate anion) in concentrated (-1OM) sodium methoxide in methanol. It is likely that the excited anion can form (and therefore fluoresce) only in very concentrated basic solutions in methanol, much the same as in concentrated aqueous NaOH solutions. Thus the excited state anion-zwitterion thermodynamics are also in evidence here. The lifetimes of the 8-quinolinolium cation and 8-quinolinolate anion, in their lowest excited singlet states, should be considerably longer than that of the neutral species. Goldman and Wehry report the quantum yield of fluorescence of the 8-quinolinolium cation in concentrated H2SOato be 0.31. A rough estimate of the quantum yield of fluorescence of the 8-quinolinolate anion in 14.6M NaOH, by the method of Parker and Rees (8) yields a value of 0.25. NOW4, = k f ~ f where 4, is the quantum yield of fluorescence, k , is the molecular probability of fluorescence (the reciprocal of the natural lifetimes of the excited state), and T~ is the radiative lifetime of the excited state. Values of k , can be approx~~
(7) M. Paul and F. Long, Chem. Rec., 57, 1 (1957). 286
(8) C. A. Parker and W. T. Rees, Analyst, 85,587 (1960).
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imated for the cation and anion from the integrated, long wavelength, absorption band intensities (9) and were computed to be 1.7 X 107 sec-1 and 2.5 X lo7 sec-1 for cation and anion, respectively. These results yield values of ~f of 2 X and 1 x lo-* sec for cation and anion, respectively. This is adequate time for excited state prototropism to be kinetically feasible. Finally, with regard to 8-quinolinolium ion fluorescence in water, in the 8-methoxyquinolinium ion, which is similar to 8-quinolinol electronically, but differs in that dissociation of the phenolic group is impossible, the characteristic green quinolinolium fluorescence prevails even in dilute acid solution and is quenched by base giving a titration curve very similar to that of the ground state cation in water (1). This indicates that dissociation occurs from the ring nitrogen and that the quenching is static. This is not inconsistent with the enhancement of 8-quinolinolium acidity in the excited state (9) N. Turro, “Molecular Photochemistry,” Benjamin, New York, 1965, p 24.
with the excited zwitterion as the conjugate base since the 8-methoxyquinolinium ion cannot dissociate to form a zwitterion. In conclusion, while hydrogen bonding is probably of considerable importance in determining the quantum yields of the various species derived from 8-quinolinol in different solvents, the fluorescence phenomena observed in solutions of differing acidities and basicities are primarily due to prototropic equilibria in the lowest excited singlet states of the various prototropic species derived from 8-quinolinol. STEPHEN G. SCHULMAN Department of Pharmaceutical Chemistry College of Pharmacy University of Florida Gainesville, Fla. 32601 RECEIVED for review September 8,1970. Accepted December 3, 1970.
Role of Solvating Agents in Promoting Ion Pair Extraction SR: In an earlier publication ( I ) , we pointed out the role of specific solvation in promoting ion pair extraction of organic ions from aqueous solutions. In the past, ion pair extraction has been related to molecular weight, degree of branching on an aliphatic chain, and dielectric constant of the solvent as well as solvation (2-4). Divatia and Biles (2) suggested that one could use the dielectric constant of the organic phase as a guide in choosing a suitable solvent for extraction. In a later article, Hull and Biles (3) analyzed the distribution of some amine salts of aromatic sulfonic acid dyes between water and less polar solvents in terms of specific solvation and concluded that the formation of solvated species is more important in the distribution than are parameters such as dielectric constant. In a recent communication, Freiser (5) has questioned interpretation of ion pair extraction by either correlation with the dielectric constant values of the solvents or by attributing promotion of ion pair extraction of organic ions to solvate formation. Rather, he suggests that such data can be better explained using the theory of regular solutions (6). In general, no single, current theory can adequately describe all extraction systems. In some cases, specific solvation of the species extracted clearly i s the dominant $actor affecting the degree of extraction. In other cases, specific solvation appears to be negligible and the extraction can be attributed primarily to van der Waals types of interactions. Regular solution theory was developed for those systems falling in the latter category. Ion pairs are definitely polar species, and thus one is inclined to be wary of interpretations based on the assumptions of regular solution theory, In our earlier report ( I ) we attributed, for example, the linear relationship between the log of the extraction constant (1) T. Higuchi, A . Michaelis, T. Tan, and A. Hurwitz, ANAL, CHEM., 39, 974 (1967). (2) G. J. Divatia and J. A. Biles,J. Pharm. Sci.,50,916 (1961). (3) L. R . Hull and J. A. Biles, ibid., 53, 869 (1964). (4) P. Mukerjee, ANAL. CHEM., 28, 870 (1956). (5) H. Freiser, ibid., 41, 1354 (1969). (6) J. H. Hildebrand and R. L. Scott, “Solubility of Nonelectrolytes,” 3rd ed., Reinhold, New York, N.Y., 1950.
of dextromethorphan hydroiodide against the log of the chloroform concentration in cyclohexane (used as the extracting solvent) to specific solvation of the ion pair by chloroform which enhanced the extraction. By plotting the log of the extraction constant against the average solubility parameter, 6, of the solvent mixture, Freiser ( 5 ) showed that one also obtains a straight line for the data mentioned above. This led him to suggest that rather than invoking specific solvation these data could be better explained on the basis “that the solubility of a solute increases with decreasing difference between its solubility parameter, 6, and that of the solvent.” He also states that “postulation of specific ‘chemical’ interaction between solvent and solutes of the ion pair type may be superfluous in many cases and that use of other solvents having solubility parameter values even greater than 10 may well result in still higher extraction constants with such solutes.” We now present some previously unpublished data which support an interpretation based on specific solvation rather than one based on the concept that extraction is always greater when the solubility parameters of the solvent and solute are more alike. The logarithms of experimentally determined extraction constants for the dextromethorphan-HBr pair in chloroformcyclohexane, cyclohexanone-cyclohexane, and nitrobenzenecyclohexane solvent mixtures are shown plotted as a function of the log of the molar concentration of the more polar solvent species, the log of the average dielectric constant ( E ) , and average solubility parameter (6) of the mixed solvent in Figures 1, 2, and 3, respectively. It will be noted that except for the cases of cyclohexanone in Figure 3, straight line correlation is approximately observed. In the case of the cyclohexanone-cyclohexane system, points are plotted only for very dilute and concentrated solutions. At intermediate cyclohexanone concentrations, emulsification results. Since straight line correlation is observed with the dilute and concentrated solutions, it seems reasonable to assume that such is also the case at the intermediate concentrations. The fact that straight lines are generally obtained as a function of 6 as well as the log of the concentration in the examples
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