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Cresyl violet chemistry and photophysics in various solvents and micelles
.7. Photochem. Photobiol. A: Chem., 64 (1992) 343-358
343
Cresyl violet chemistry and photophysics in various solvents and micelles Stefan
J. Isak and Edward
Depa~ent
M. Eyring
of Chemirtry, Univeni~ of Utuh, SaIt Lake City, UT 84112 (USA)
(Received August 28, 1991; accepted January 16, 1992)
Abstract The chemical and photophysical aqueous micellar and organic
behavior of cresyl violet perchlorate was studied in aqueous, solutions. In aqueous micellar solutions the fluorescence
quantum yield of cresyl violet increases in sodium dodecylsulfate (SDS), whereas a dramatic decrease is observed in Triton X-100 (TX). The chemical and photophysical behavior of cresyl violet is strongly dependent on hydrogen-bond, electrostatic and acid-base interactions in solution, as well as un its hydrophobicity.
1. Inh-oduction used as a laser dye. It is a cationic Cresyl violet (CV’) (Fig. 1) is commonly chromophore which absorbs red light and has a fairly high fluorescence quantum yield &. The highest & for CV+ is observed in alcohols. A value for & of CV+ in dilute methanolic solutions (less than 10e6 M) of approximately 0.70 has recently been determined [l]. CV+ has also been applied in organic chemical sensors [2], solar energy conversion and energy storage [3], as a photosensitizer in energy and electron transfer reactions [4] and as a potential photosensitizer in photodynamic therapy [5]. Thus a clearer understanding of the chemistry and photophysics of CV+ would be of great interest.
Fig. 1. The two major resonance structures of cresyl violet perchlorate.
lOlO-6030/92/$5.00
0 1992 - Elsevier Sequoia. All rights reserved
344 CV+ is an oxazine-type dye. This class generally has properties similar to the related xanthene and rhodamine dyes [6-S]. CV+ possesses an insignificant triplet population on excitation and is found to be quite photochemically stable [6, 91. CV+ is a highly rigid, planar molecule with the amine hydrogens lying in the plane of the molecule [6]. The amino groups are relatively acidic [6, 81, and CV+ and related dyes have been observed to be deprotonated in basic solution 15, 8, 10-K?]. The major sources of non-radiative decay arise from the small energy difference between the ground and excited states [6] and the high frequency N-H vibrations of the amino groups [B, 131. As the energy gap between the ground and excited states decreases, these vibrations become more important in non-radiative decay [8]. Because of its structure, CVf has the potential to dimerize and does so extensively in aqueous solution where dimer formation is assisted through hydrogen bonding with water [14, 151. The hydrophobic@ of CV+ is also a contributing factor to dimer formation [14, 161. Few studies concerning CV + in particular have appeared in the literature. More often than not, CV+ has been included as part of a more general study and, as a consequence, different conclusions regarding its behavior have been reached. Our initial interest in CV+ was to use it as a fluorescence standard, and thus its fluorescence quantum yield was characterized in methanolic and aqueous solutions as a function of concentration [l]. In considering CV+ for use as a laser dye, we would like to determine its behavior in aqueous micellar solutions, since the use of aqueous micellar solutions of laser dyes has a number of advantages over other solvents [6, 171. With these considerations in mind the concentration-dependent r#+ values of CV+ in aqueous micellar solutions of three common surfactants were investigated. At the same time the absorbance and fluorescence characteristics of CV’ in a variety of solvents were investigated to provide a more comprehensive understanding of the chemistry and photophysics of CV+.
2. Experimental details 2.1. Absolute fluorescence quantum yields of cresyl violet in methanol Absolute values of & for CVf in methanol were determined using photothermal beam deflection photoacoustic spectroscopy (PBDPAS) [l]. The theory of this technique has been treated by both Sullivan and Tam [l&l and Terazima and Azumi [19]. The experimental set-up is shown in Fig. 2. The pump laser was a Quanta Ray DCR-2 Nd:YAG laser from which the frequency-doubled h = 532 nm output was used to pump a Lambda-Physik FL2002 dye laser. The laser dye employed was rhodamine 640 perchlorate. An excitation wavelength of 610 nm from the dye laser was used as the pump beam. The pulse repetition rate was 10 Hz and the energy of the pulse was varied between 0.1 and 1.0 CLJ.A 3 mW diode laser (A=780 run) or a 5 mW HeNe laser was used to supply the probe beam. Excitation and probe energies were adjusted using neutral density filters. The probe beam signal was detected with a photomultiplier tube (Hamamatsu R928) wired for fast response [20]. The photoacoustic signals were averaged on a LeCroy 9400 oscilloscope and stored on disk by an IBM PC compatible computer. Sample cells were made of quartz with an optical path length of 1 cm. Measurements were made at an ambient temperature of 24.0&0.2 “C.
345
photomultiplier tube
probe
laser
m X.YBZ
platform
splitter
X.Y.2
platform
Fig. 2. Experimental arrangement for the thermal lens measurements. The output from an Nd:YAG-laser-pumped dye laser is directed through a diaphragm. An HeNe or diode laser probe beam is made collinear with the pump beam using a beam splitter. The positions of the pump and probe beams as well as their focusing can be additionally controlled with a lens. The pump and probe beams are then directed approximately 2.5 m beyond the cell using a series of prisms. The probe beam is then cut approximately in half with a razor edge and focused onto a photomultiplier tube (PMT) using another lens. A filter is placed in front of the PMT to filter out any light from the pump beam. The output from the PMT is displayed on an oscilloscope where the waveform may be plotted or stored on disk by a computer. (Reprinted from S. J. Isak, S. J. Komorowski, C. N. Merrow, P. E. Poston and E. M. Eyring, Appl. Spectrosc., 43 (1989) 419. Copyright 1989, Society for Applied Spectroscopy, Frederick, MD.) 2.2.
Absorbance and fluorescence measurements
A Perkin-Elmer MPF-66 fluorescence spectrophotometer was used for the fluorescence measurements. Emission spectra were corrected using a totally reflecting surface supplied by the manufacturer. Absorbance measurements were made at an ambient temperature of 24.0 I,!I0.2 “C using a Perkin-Elmer Lambda 9 spectrophotometer. 2.3. Refractive index measurements The refractive indices of the aqueous micellar solutions at ambient temperature were determined with a Bausch & L-omb Abbe-3L refractometer. The results are included in Table 1 (see Section 3). 2.4. Equations Equation (1) [21-281 was used to determine the absolute fluorescence quantum yield of CV+ in methanol from the PBDPAS measurements (1)
346
where hF is the mean fluorescence wavelength determined from corrected emission spectra. A, is the excitation wavelength, H,, is the non-radiative emission intensity of emission intensity of a non-luminescent the sample and H,* is the non-radiative reference. The non-luminescent reference sample used was CV* in KI-saturated methanol. KI completely quenches CV’ fluorescence and produces no changes in the absorbance spectra. The fluorescence quantum yields of CV+ in aqueous and aqueous micellar solutions were determined relative to the absolute & values of CV+ in methanol according to the following relationship [29, 301
where &,ref is the fluorescence quantum yield of a reference compound (in this case the absolute & of CV’ in methanol),A is the absorbance, nD is the index of refraction of the solvent and a is the area under the fluorescence peak. Equation (2) is valid as long as Beer’s law is obeyed and there is no inner filter effect [31-371 on the fluorescence. For CV+, eqn. (2) is only valid for concentrations below approximately lO-‘j M [l], Above this concentration of CV+, dimerization and/or inner filter effects cause problems. To compensate, solutions of matched absorbances were used to determine & values at CV+ concentrations above 10m6 M. 2.5. Maserials EU (Mallinckrodt) was recrystallized from water and subsequently dried and stored in a desiccator under vacuum. Rhodamine 640 perchlorate and cresyl violet perchlorate from Exciton, sodium dodecylsulfate and hexadecyltrimethylammonium bromide from Sigma, Triton X-100, N,N-dimethylformamide (high performance liquid chromatography (HPLC) grade), methylsulfoxide (HPLC grade), triethylene glycol dimethyl ether and tetraejhylene glycol from Aldrich, acetone (PHOTREX; spectrophotometric grade) and ethyl ether (anhydrous) from J. T. Baker and hexane (Omni-Solv; spectrophotometric grade) and methanol (Omni-Solv; spectrophotometric grade) from EM Science were all used as received. Ultrapure water was used from a Corning MP-190 water purification system connected to an MP-1 water still.
3. Results and discussion 3.1.
Solubilization
and
ground state characteristics of cresyl violet
Absorbance maxima, mean fluorescence wavelengths and fluorescence quantum yields for dilute aqueous, aqueous micellar and organic solutions of CV’ are listed in Table 1. A number of solvent electro-optical properties, as well as the apparent solubility of CV+ in these solvents, are also listed. The solvents are presented in the order in which the’absorbance maximum of CVc is observed to increase (i.e. shift to longer wavelengths). A brief glance at Table 1 suggests little, if any, correlation between the solubility or absorbance maxima of CVf and any electro-optical property of the solvents listed. However, closer inspection of the data yields a definite correlation between the solubility and absorbance maxima and a number of solute and solvent properties. These include the polar and electrostatic solute-solvent interaction, the hydrophobic@ of CV+, the
(CO.40)
617 620 621 624 628 626 627 629 629
589 (462) 593 (472) 594 598 599 601 603 607 607
‘Ref. 38. b20 “C. ‘Ref. 39. dRef. 40.
(0.40-0.67) (0.40-0.67) 0.67 0.52 (<0.40)
(=S94)
(=595) 623 624
480 (603) 482 585 589
0.22 (x0.40) ( < 0.40)
(<0.40) (<0.40) 0.40 0.40
(<0.40)
(5589)
473 (604)
Triethylene glycol dimethyl ether Ethyl ether N,N-Dimethylformamide Water Hexadecyltrimethylammonium bromide (10 mM) Acetonitrile Acetone Methanol Sodium dodecykulfate (70 mM) Tetraethylene glycol-water (50:50) Triethylene glycol dimethyl ether-water (50:50) Triton X-100 (30 mM) Tetraethylene glycol Methylsulfoxide Hexane
4F
Mean fluorescence wavelength (nn.0
Maximum absorbance wavelength (nm)
Solvent
High High High Insoluble
High High High High
Very low High ~10-~ M
Low (-~10-~ M)
Solubiliry
1.96d 0.294’
46.6d 1.89”
1.3360f 0.0004 1.458” 1.476” 3.96” 1.372” 0
1.342” 3.92” 1.357’ 2.88’ 1.326’ 1.70a 1.3355f 0.0006 0.34Y 36.2d 0.316” 20.7” 0.547” 32.6’
Dipole moment (debye)
1.352’ 1.15’ 1.427’ 3.82’ 1.333” 1.85’ 1.3336f 0.0005
$5 “C)
0.222’ 4.34a 0.796 36.7” 0.890” 78.5”
y2.S“C) ;25 “C) (cP)
Physical and spectroscopic properties of ditute (less than 10e6 M) solutions of cresyl violet in a number of solvents
TABLE 1
4
w
348
acid-base equilibrium of CV+ in solution and the hydrogen-bonding characteristics of the solvent. The importance of hydrogen-bonding interactions of the solvent with CV+ and related dyes has previously been recognized [41-44]. Of the pure solvents listed in Table 1, three can be considered to be amphiprotic, i.e. capable of acting as either hydrogen-bond donors (HBDs) or hydrogen-bond acceptors (HBAs) [45, 461. These are water, methanol (MeOH) and tetraethylene glycol (TEG). The remaining pure solvents, with the exception of hexane, may all be considered to be HBAs since none of them has a hydrogen available for proton-donating interactions. Hexane is nonhydrogen bonding (NHB). Since water is a component in all of the mixed solvent systems, the mixed solvents may also be considered to be amphiprotic. CV+ has a number of sites that may be involved in a variety of interactions. These are indicated in Fig. 3. The specific hydrogen-bonding interactions may be described as follows. The aromatic amine group may be involved in two types of interaction. The amine hydrogens can hydrogen bond with an HBA, e.g. with the oxygen atom in water or methylsufioxide (DMSO). The amine nitrogen may hydrogen bond with an HBD, e.g. with the hydrogen in water or MeOH. The heterocyclic nitrogen and oxygen may hydrogen bond with the hydrogen from an HBD. Finally, the hydrogen atoms of the protonated amine may interact with an HBA. Thus, in terms of hydrogen bonding, we might expect water, MeOH and TEG to provide the greatest solvating power. On the basis of ground state stabilization’we might expect the absorbance maxima of CV’ in these three solvents to be blue shifted with respect to the remaining solvents. This is because a red shift in the absorbance maximum with a decrease in hydrogen-bonding interaction is expected for CV+ 147-503. However, as is evident from Table 1, this is not the case. Also inconsistent in terms of hydrogen bonding is the low solubility of CV+ in water. Thus the solubility and stabilization of CV’ do not arise solely from hydrogen-bonding interactions of CV’ with the solvent. The limited solubility of CV” in water may be explained by the somewhat hydrophobic characteristics of CV+. Although the benz[a]anthracene skeleton of CV+ has been substituted by a number of polar and hydrophilic substituents, it still possesses considerable hydrocarbon character. The polar nature and electrostatic charge of C!V+ also play an important roIe in its solubilization and stabilization. Thus amphiprotic solvents provide extensive solubilization and stabilization for both CV+ and the perchlorate (ClO,-) anion. In the remaining solvents, which are all HBAs with the exception of hexane, extensive
H2y&; 8 H I
/
Ii
CIOF d
0
!
“\,/”
l-i2’---
.
/
0
\
i
“\,/”
H R
s
Hc;\ R
R
Fig. 3. Potential sites for hydrogen-bonding, electrostatic violet (see text for a discussion of these interactions).
and acid-base interactionsin cresyl
349
solubilization and stabilization of CV +, but not C104-, can be expected. The poorly solvated ClO.+- is thus potentially quite reactive. The degree of electrostatic interaction between solute and solvent should correlate fairly well with the dielectric constant or dipole moment of the solvent, the solvent with the largest value of the dielectric constant or dipole moment providing the most solubilizing and stabilizing power, Such electrostatic considerations may explain why the absorbance maximum of CV+ in acetonitrile (CE&CN) is blue shifted with respect to that in acetone, but do not explain why the absorbance maximum of CV+ in NJV-dimethylformamide (DMF) is significantly (greater than 100 nm) blue shifted with respect to that in CHsCN, which has a comparable dielectric constant, or to that in DMSO, which has a comparable dipole moment. The large blue shift of the absorbance maximum of CV4 in triethylene glycol dimethyl ether (triglyme), ethyl ether and DMF is owing to the deprotonation of CVc. Thus an acid-base equilibrium exists in solution between CV’ and the solvent. If the solvent is sufficiently basic and capable of accepting a proton, a transfer of one of the protonated amine hydrogens to the solvent may take place. Such a transfer would significantly stabilize CV+ and consequently result in a considerable blue shift of the absorbance maximum. The ethers and DMF are basic solvents [51-531 capable of accepting a proton. In the case of DMF, considerable resonance stabilization of the protonated form may be expected. Similarly, large blue shifts in the absorbance maximum, owing to the deprotonation of CVf, have been observed in a number of solvents, including DMF [6, 10, 12, 541. The addition of water or a small amount of acid to the deprotonated solutions of CVf restores the protonated form as illustrated by a red shift in the absorbance maximum back to approximately 600 nm. This can be clearly seen to occur with the triglyme and 50~50 water-triglyme solutions as indicated in Table 1. However, in at least one report [4], complex formation between CV+ and the basic solvent is suggested as opposed to the acid-base interpretation of the data discussed here. In some cases, the absorbance bands of both the protonated and deprotonated forms of cresyl violet can be observed. When this occurs one species is predominant. For the solvents in which this is observed the absorbance maximum of the minority form is indicated in parentheses next to the absorbance maximum of the predominant form. Some workers [13, 41,551 have attributed the spectroscopic characteristics of CV+ to interactions of the solvent solely with the heterocyclic nitrogen and/or oxygen in CV+. This seems unlikely based on the present discussion and other reports [6, 10, 12, 541. The solubility and absorbance data in Table 1 can thus be explained in terms of the hydrogen-bonding interactions between CV+ and the solvent, the somewhat interactions between CV+ and the hydrophobic characteristics of C!V+, electrostatic solvent and the acid-base equilibrium between CV+ and the solvent. While other contributions to the observed spectroscopic properties of CV+, such as the effect of the solvent refractive index on the absorbance maximum, are, no doubt, superimposed, they can be expected to be overshadowed by the contributions made by the acid-base, hydrogen-bonding and electrostatic properties summarized above [56]. The important solvent-solute interactions summarized above for CV+ facilitate an understanding of the difference in CV+ behavior in various solvents, even though it may be difficult to judge the relative importance of each in the overall solubilization and stabilization of CV+. For example, it should now be clear why CV+ is not soluble in hexane. First, hexane is non-polar with a very low dielectric constant. Thus little
350
Secondly, no hydrogen-bonding if any electrostatic interaction with CV + is expected. interactions are expected between CV+ and hexane, since hexane has no hydrogenbonding capacity, either as an HBA or HBD. Similar observations have been made for CV+ in other non-polar solvents [7, 415 In a comparison of the characteristics of CV+ in water, CH&!N, acetone and MeOH, the limited soIubility in water has already been attributed to the somewhat hydrophobic nature of CV+. The increase in the absorbance maximum of CV+ in the order water < CH&N < acetone < MeOH can be explained in terms of hydrogenbonding and electrostatic interactions between CVf and the solvent. Water and MeOH would provide maximal hydrogen-bonding interactions with CV+ serving as both HBAs and HBDs. The blue shift in water with respect to the other solvents is attributable to the strong hydrogen bonds formed and its high dielectric constant. The blue shift of CV’ in CI-&CN with respect to MeOH may be attributed to the fact that CH,CN is considerably more polar with a somewhat larger dielectric constant than MeOH. This results in a greater stabilization of CV+ in CH,CN than in MeOH. The absorbance maximum of CV+ in acetone is virtually the same as that in MeOH. While acetone is more polar than MeOH, it has a lower dielectric constant. The absorbance spectrum of CVf in acetone indicates the presence of both the protonated and deprotonated forms. As previously discussed, the deprotonation of CV+ results in considerable stabilization. These factors apparently contribute to the stabilization of CV’ to provide a ground state which is similar in energy (if not slightly lower than) in acetone to that in methanol. 3.2.
Excited state charactetitics
of cresyl violet
In all solvents, with the exception of DMF and the ethers, there is a significant overlap between the absorbance and fluorescence spectra of CV+. There is also a small Stokes’ shift. The largest Stokes’ shifts for CV’ are observed in water and aqueous hexadecyltrimethylammonium bromide (CTAB) solutions. As the absorbance and fluorescence spectra are progressively red shifted, a general decrease in the Stokes’ shift is observed, with the TEG and DMSO solutions exhibiting the smallest Stokes’ shifts. The range of the shift in the absorbance maximum is nearly twice that observed for the mean fluorescence wavelength hF. The absorbance spectra reveal a shoutder that is blue shifted with respect to the absorbance maximum. The fluorescence spectra are fairly good mirror images of the corresponding absorbance spectra. The fluorescence spectra were obtained by exciting the sample at 610 nm. The situation is quite different in DMF and the ethers. In these solvents deprotonated cresyl violet (CV) displays a very large Stokes’ shift and the absorbance and fluorescence spectra differ markedly. The absorbance spectra are centered around approximately 480 nm with fairly symmetrical peaks. The blue-shifted shoulders observed for CV+ are no longer present. The fluorescence spectra are quite broad, approximately 85 nm at full width at half-maximum (FWHM) compared with approximately 30 nm at FWHM for CV+. As a consequence, the fluorescence spectrum of CV overlaps the region of fluorescence of CV+ to a considerable degree. Owing to the lack of absorbance at 610 nm the fluorescence spectra of CV in these solvents were obtained by excitation at 490 nm. The fluorescence spectra display some structure that appears to be resolvable into four separate fluorescence peaks. Such an observation is unusual. The fact that excitation-wavelength-dependent emission spectra have been observed for similar deprotonated oxazine dyes [lo] suggests that the observed fluorescence spectra may depend on the excitation wavelength used. It is also curious that the observed fluorescence spectra are virtually the same in the three different solvents.
351
For the solvents in which b of C!V+ was determined, the highest value was observed in MeOH and the lowest in aqueous Triton X-100 (TX) solutions. The & values are listed in Table 1. Qualitative assessment of the J#+ values of C!V+ in acetone and CH3CN indicates that they are similar to that in MeOH and probably not less than the value in water. In the remaining solvents, the fluorescence is quenched with respect to that in water. The & values of CV in DMF and the ethers were more difficult to ascertain because of significant differences in absorbance. However, a & value less than that in water is expected [4, 10, 54, 57, 58). The estimated values of & for CV+ and CV are listed in parentheses in Table 1 together with the measured vaiues. The relatively small Stokes’ shift and large overlap of the absorbance and fluorescence spectra indicate small configurational changes between the ground and first excited states of CV+ [59]. This is confirmed by the good mirror symmetry of the absorbance and fluorescence spectra, indicating negligible reorganization effects and small geometry changes between the ground and first excited states [60]. The larger shifts observed for the absorbance maxima compared with I\Findicate that the ground state may be more polar and more solvent stabilized than the first excited state. This is reasonable if the large blue shift in absorbance on going from CV+ to CV is indicative of a general increase in the basicity of CV’ on excitation [SE&61-J. If this is the case, the importance of electrostatic and acid-base interactions of the solvent with the protonated amine group probably decreases on excitation. Such an interpretation also explains the larger shifts observed for the absorbance maxima compared with & and the general decrease in the Stokes’ shift as the absorbance and fluorescence spectra are progressively red shifted. For example, in a comparison of the data for water and DMSO, we would expect water to provide the most significant hydrogen-bond interactions and also significantly stronger electrostatic interactions because of the high dielectric constant. This results in a blue shift of the absorbance spectrum in water relative to that in DMSO. If, on excitation, some of the ground state interactions become less important, a larger energy difference is expected for water than for DMSO and is expressed in the larger Stokes’ shift. A similar observation in rhodamine 6G solutions has been explained as being due to orientational reorganization of nearby polar groups f17]. Such reasoning is also consistent with the conclusions reached regarding solvent effects on merocyanine dyes [49]. The absorbance spectrum of CV+ displays a blue-shifted shoulder. This shoulder disappears in the case where CV is formed. This blue-shifted shoulder has been ascribed to steric hindrance produced between the hydrogen of the amine group and the hydrogen of the phenyl group located nearby [63 or to two different oscillator groups within CV+ IlO] which are nearly degenerate. On deprotonation of CV+ the major resonant forms pictured in Fig. 1 will lose their resonance and the corresponding isomers will be formed. However, one of these isomers will still have the amine and phenyl groups cis to one another so that steric hindrance may still occur. There is no blue shoulder in the absorbance spectrum and the argument favoring similar oscillator groups in CV+ IS - more persuasive. As a brief aside, there appears to be some debate over whether the excitation and/or relaxation of CV+ may involve a charge transfer (CT) process [4, 58, 62, 651. A detailed discussion is beyond the scope of this report and may be found elsewhere [48,50,58, 661. However, it should be clear that the unprotonated amine or heterocyclic amine of CV+ may act as an electron donor and the protonated amine group may act as an electron acceptor. Such CT processes are expected to quench signscantly any fluorescence [50,58,62+4,66] and are not consistent with the significant fluorescence
352
observed. Thus it appears that a CT process is not operative. This lack of a CT process may be attributed to the greater rigidity of the aromatic amine groups of CV’ compared with the more bulky alkyl amino groups found in other oxazine and similar dyes for which CT processes have been observed 18, 64, 67, 681. CT processes do appear to be operative in the case of CV. These CT processes are a consequence of the acid-base equilibrium that can exist between CV+ and CV in solution. Two CT mechanisms have been proposed. The first mechanism involves the formation of a CT complex [4]_ CV is characterized by a significant blue shift in the absorbance maximum with respect to CV’ and its fluorescence is greatly diminished, displays a large Stokes’ shift and overlaps, to some extent, the region where CV+ is observed to fluoresce. In ref. 4, the large blue shift of the absorbance maximum of CV+ in amines is attributed to the formation of a CVf-amine complex. The large Stokes’ shift of the fluorescence spectrum of the complex and its coincidence with the fluorescence spectrum of uncomplexed CV+ are attributed to the similar fluorescence characteristics of uncomplexed and complexed CV+. Some of the complexed CV+ has been observed to dissociate on excitation. This dissociation is suggested to be owing to a smaller K, value in the excited state which results in a weakening of the CT interaction in the CV+-amine complex. Dissociation of the complex on excitation is indicative of a more basic excited state and thus of an increase in the pK, value of CVf on excitation as has previously been suggested. The other CT mechanism involves the transfer of a proton between CV+ and the solvent [lo, 11, 54, 651 and has already been discussed to some extent. In basic proton-accepting solvents a proton may be transferred from the protonated amine group of CV’ to the solvent. The stabilization obtained on formation of CV results in a considerable blue shift of the absorbance spectrum. Evidence suggesting that CV’ becomes more basic on excitation was discussed in the preceding paragraph. pK, changes between the ground and first excited states can be very large, with differences of 5-10 pK, units [SO, 58, 691 not uncommon. Thus, on excitation, it is possible for CV to become basic enough to be reprotonated to some extent, if not fully. Such a proton transfer process is expected to quench considerably any fluorescence [50, 58, 62-64, 661. We might also expect the observed fluorescence spectrum of CV to be similar to that of CV’. Both expectations are consistent with the data. The fluorescence is quenched in all cases for CV, there is considerable overlap of the fluorescence spectrum of CV with that of CV+, and the large Stokes’ shifts for CV indicate a ground state more stable than that of CV+, but a similar excited state. Proton transfer may also explain the fluorescence spectra observed for CV. The fluorescence spectra are broad and can be resolved into four fairly equivalent peaks. Two of these peaks are identical in intensity as are the other two. The differences in intensity are less than 10%. The slightly less intense peaks are blue shifted relative to the slightly more intense peaks. Thus it seems that there are two pairs of peaks. How they may be related is not clear. From Table 1 it can also be seen that the & value of the two slightly more intense peaks is always the same, while the hF value of the slightly less intense peaks varies with the solvent. There are two major resonance forms of CV+ (see Fig. 1) and, on deprotonation, this resonance is lost and the two corresponding isomers are formed. If steric hindrance is not a factor, a 50:50 mixture of the isomers may be expected. If the isomers of CV are fluorescent, as is CV+, it is possible to have four fluorescing species in solution at the same time. The observability and intensity of each fluorescent species are clearly dependent on the kinetics of all the excitation and de-excitation processes involved. The kinetic considerations involved have been considered elsewhere 150, 581.
353 The two isomers of CV are nearly, but not exactly, the same, just as the two different oscillator groups [lo] observable in the absorbance and fluorescence spectra of CV+ are nearly, but not exactly, the same (as discussed above). Such a difference is most clearly manifested in the excited state. The stabilizing power of the solvent may be similar for both the ground and excited states of CV. In this case we would expect to observe solvent-dependent IF values as is the case for the slightly lower intensity fluorescence peaks of CV. On the other hand, if CV is reprotonated on excitation, solvent interactions may be reduced and become less important as in the case of CV+. Such reasoning is supported by the generally solvent-independent AF values observed for CV+. On reprotonation of the CV isomers, resonance will be re-established. Although individual absorbance and fluorescence peaks of the resonance forms were not observed for CV+, it is apparent that some distinction is possible as illustrated by the shoulder observed in the absorbance and fluorescence spectra of CV+ . Such small energy differences may be manifested in the fluorescence spectrum of CV. The stabilization provided by DMF and the ethers may easily allow for a distinction between the two species. This suggestion is supported by a trend in the data indicating a decrease in the peak width (FWHM) of the fluorescence with an increase in the blue shift of the fluorescence spectrum, so that in an extremely stabilized environment, the two peaks may be resolved. Finally, the variation in 6 of CV + in different solvents merits discussion. While proton transfer appears to be predominantly responsible for the quenching of the fluorescence of CV, this is not the case for CV’. For CV+, hydrogen-bonding and electrostatic interactions can be expected to be important as well as any process that may contribute to an increase in the non-radiative decay rate. For example, the following factors all provide possible explanations for the greater quenching of CV + in DMSO than in MeOH. First, as a general rule, the rate of non-radiative decay increases with decreasing energy difference between the ground and excited states [68]. Secondly, and perhaps more likely, there is quenching of the fluorescence by the C104- counter-ion. From Table 1 it is apparent that little stabilization of CV+ is obtained through its interaction with DMSO. At the same time, owing to the nature of polar aprotic solvents, the C104counter-ion is poorly solvated and rendered potentially quite reactive. Thus the poor solubilization and stabilization will lead to a more tightly associated CV’-C104interaction in which the ClO,ion may quench the fluorescence of CV +. Finally, it is possible that the sulfur atom of DMSO may serve as an effective fluorescence quencher. The fluorescence quenching ability of sulfur, primarily through increased triplet yields, has been reported [5, 61. For example, the fluorescence of thiazines is quenched with respect to oxazines, because of the presence of a sulfur atom in the molecule. While most instances appear to involve intramolecular sulfur quenching, it is possible that an efficient interaction of the sulfonyl oxygen of DMSO with the cationic amine nitrogen will provide sufficient conjugation of DMSO with the r electronic system of CVc to result in effective quenching, where sulfur acts as the electron acceptor. The fluorescence of CV+ in water is also quenched with respect to that in MeOH for a variety of reasons. One contributing factor is the formation of non-fluorescent dimers in water as noted above. The importance of N-H vibrations in the non-radiative deactivation of CV+ has also been mentioned. Additional vibrational energy loss may occur for C!V+ through 0-Hvibrations with hydrogen-bonded water. The CV+-hydrogen bonds are expected to be both strong and efficient. Such additional energy losses have been suggested by Fiirster [70] and observed in the case of rhodamine B 1681.
354
3.3. Interaction of cresyl violet with aqueous micellar solutions The variation in & of CV* in water and aqueous micellar solutions of CTAB, sodium dodecylsulfate (SDS) and TX is presented in Fig. 4. The & values for the dilute micellar solutions of CV+ are listed in Table 1. From Fig. 4 it can be seen that the concentration-dependent & behavior of CV+ in CTAB is virtually identical to that in water. In SDS micellar solutions, an enhancement in the fluorescence of approximately 25% is observed, as opposed to a decrease in & of approximately 50% in micellar TX solutions. There is an increase in the observed absorbance maximum and & in the order water
concentration
Fig. 4. Concentration-dependent aqueous CXAB
[Id]
fluorescence
quantum yields of cresyl violet in aqueous (v),
(Cl), aqueous SDS (A) and aqueous TX
(0)
solutions.
355 the interaction between C!V+ and CI’AB which takes place in the ground state is diminished in the excited state. Interaction of CV+ with the Gouy-Chapman layer of the micelle may account for this. The Gouy-Chapman layer consists of a high density of counter-ions of the polar CX4B head groups, in other words Br- and C104-. On excitation, CV+ becomes more basic and interactions with such negatively charged species will diminish. It is interesting to note that no quenching of the fluorescence of CV* is observed as a consequence of its potential for interaction with Br-, a fairly effective quencher. Perhaps the interaction of CV+ with the Gouy-Chapman layer is inconsequential or CVc, Br- and C104are effectively shielded from one another as a consequence of the intense soIvating power of water. SDS is an alkyl sulfate anionic detergent. The negatively charged sulfate head group is counterbalanced by sodium Na +. There is no evidence for CV+ dimerization in the SDS micellar solutions. The polar micelle-water interface of SDS [7&82] and the complementary electrostatic charges of CV+ and SDS both favor a strong interaction between the two. In addition to electrostatic interactions, some hydrophobic interaction of the hydrocarbon-like portions of CV+ may be expected with the long hydrocarbon tails of SDS. Consequently, it is possible that CV+ may be localized at the micelle-water interface. In general, solute-micelle interactions become stronger as the solute becomes more hydrophobic [47, 83-861. The hydrophobic@ of CV+ has already been discussed. The hydrophobicity of CV+, the observed loss of dimerization of CV+ in SDS and the significant shift in the absorbance maximum of CV+ in SDS with respect to that in water all indicate a strong interaction of CV+ with SDS. Hydrogen-bonding interactions between CVf and water will be diminished on interaction of CV+ with the micelle. This is reflected by the red shift in the absorbance maximum of CV+ in SDS with respect to that in water, consistent with the observation that a decrease in hydrogenbonding interaction is expected to produce a red shift in the absorbance spectrum of CV+_ The lack of a significant difference in the & values between C!V+ in water and SDS may be explained as previously for CTAB. The & value of CV+ in SDS, while enhanced with respect to that in water, is still quenched with respect to the value in MeOH. So, although the fluorescence may be enhanced owing to the elimination of dimers and perhaps other factors, other quenching processes are still active or may be generated as a consequence of the interaction_ Two possibilities for quenching include strong hydrogen-bond interactions between CV* and water in the micellar solutions and energy loss via a strong electrostatic interaction between CV+ and the polar sulfate head group of SDS. TX is a non-ionic detergent and consists of a bulky alkyl phenyl head group and a long polyoxyethylene (POE) chain which terminates in a hydroxyl (-OH) group. This renders the POE chain amphiprotic in nature with the ether oxygens capable of serving as HBAs and the -OH group as an HBD or HBA. The micelles of TX differ from those of CT’AB and SDS. The long POE chains of TX are polar and thus extend out into the surrounding aqueous environment, whereas the long hydrocarbon taits of CTAE3 and SDS are located within the interior of the micelle. The long POE chains of TX have been found to be quite flexible and may interact with one another [87-901. Evidence for a strong interaction of CV+ with TX is given by the significant red shifts observed in both the absorbance maximum and & with respect to those obtained in water, the elimination of any evidence of dimerization of CV* and a decrease in the h value of CV+ with respect to that obtained in water. Reports of fluorescence quenching in micellar solutions are not common. In one such report [91], quenching was suggested to be attributable to impurities in the TX used. This possibility was
356
ruled out in our results when enhanced fluorescence was observed in other dye systems in our laboratory [9] using the same source of TX. On the basis of the observations and discussion of CV+ up to this point, an interaction of CVf with the POE chain seems quite likely. To test this hypothesis, the absorbance and fluorescence of CV+ were studied in triglyme (an HBA), TEG (amphiprotic) and SO:50 aqueous mixtures of each. The results are listed in Table 1 and a strong interaction, accompanied by a decrease in the & value of CV+, is indicated in each case. In triglyme, the ether oxygens are apparently basic enough to deprotonate CF. The addition of water, as in the 50~50 aqueous triglyme mixture, restores CV+ . However, compared with triglyme the behavior of CV+ in TX much more closely resembles that in TEG, indicating that interaction with the terminal -OH group is very important. The extent of interaction of CV’ with the ether oxygens of TEG is not clear. The observation that the values of the absorbance maximum and & of CV+ in TX lie between the values obtained in TEG and in 50:50 aqueous TBG solutions indicates the importance of water in the interactions. Clearly, a considerable amount of interaction between the POE chain and water is expected and has been observed experimentally 188, 921. Although the exact nature and extent of interaction of CV+ with the TX micelles are difficult to ascertain, it is clear that the POE chains are capable of providing hydrogen-bonding, electrostatic and acid-base interactions at least comparable with those of water in a less hydrophobic environment. As a consequence of the interaction of CV+ with the POE chains, its fluorescence is quenched, probably because of interaction with the ether oxygens. Additional channels of non-radiative decay are probably created by additional interactions between C!V+, the POE chains and water. The flexibility of the POE chains may also contribute to the efficiency of the nonradiative processes. These conclusions are consistent with strong POE chain interactions observed in other instances [87, 89, 931 and a report in which the observed decrease in the fluorescence of uranyl ions in TX micelles was attributed to an interaction with the POE chains [93].
Acknowledgments This research was funded by the Department of Energy, Office of Basic Energy Sciences. The experiments were initiated under a Biomedical Research Support Grant (USPHS No. RR07092). The authors thank J. D. Spikes for the use of his spectrophotometer and fluorometer.
References 1 S. J. Isak and E. M. Eyring, J. Phys Chem., 96 (1992) in the press. J. F. Giuliani, H. Wohltjen and N. L. Jarvis, Opt. Lett,, 8 (1983) 54. J. S. Batchelder, A. H. Zewail and T. Cole, Appl. Opt., 20 (1981) 3733. D. 1. Kreller and P. V. Kamat, J. Phys. Chem, 95 (1991) 4406. L. Cincotta, J. W. Foley and A. H. Cincetta, Photo&em. Photobid., 46 (1987) 751. K. H. Drexhage, in F. P. Schgfer (ed.), me Lasers, Springer-Verlag, New York, 1977, Chapter 4. 7 K. H. Drexhage, Laser Focus, 9 (1973) 35. 2 3 4 5 6
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K. H. Drexhage, J. Res. N&l. Bur. Stand., Sect. A, 80 (1976) 421. S. J. Isak, unpublished results, 1991. J. F. Giuliani and T. W. Barrett, Specfrosc. Left., 26 (1983) 555. E. M. Goryaeva and A. V. Shablya, J. Appl. Spectrosc., 43 (1985) 750. E. B. Zhigalova and Yu. P. Morozova, Russ. x Phys. Chem., 59 (1985) 1010. G. S. Beddard, T. Doust and G. Porter, Chew. Phys., 61 (1981) 17. L. V. Levshin and I. S. Lonskaya, Opt. Spechosc., 11 (1961) 278. Yu. P. Morozova and E. B. Zhigalova, Russ. J_ Phys. Chem., 56 (1982) 1526. R. Ramnath, V. Ramesh and V. Ramamurthy, .7. Photochem., 31 (1985) 75. J. C. Mialocq, Ph. Hebert, X. Armand, R. Bonneau and J. P. Morand, j. Photochem. PhotobioZ. A: Chern., 56 (1991) 323. B. Sullivan and A. C. Tam, J. Acoust. Sot. Am, 75 (1984) 437. M. Terazima and T. Azumi, Chem. Phys_ Lett., 276 (1991) 79. J. M. Harris, F. E. Lytle and T. C. McCain, Anal. Chem., 48 (1976) 2095. D. Magde, J. H. Brannon, T. L. Cremers and J. Olmsted, III, J. Phys. Chem., 83 (1979) 696. J. Olmsted, III, J. Phys. Chem., 83 (1979) 2581. J. B. Callis, in K. D. Mielenz, R. A. Velapoldi and R. Mavrodineanu (eds.), Stundurdizution in Spectrophutometty and Luminescence Measurements: NBS Special Publication A66, US Government Printing Office, Washington DC, 1977. W. Lahmann and H. J. Ludewig, Chem. Phys. Lett., 45 (1977) 177. I. 0. Starobogatov, Opt. Spectrosc., 42 (1977) 172. D. Cahen, H. Garty and R. S. Becker, J. Whys. Chem., 84 (1980) 3384. M. L. Lesiecki and J. M. Drake, AppZ. Opt., 21 (1982) 557. G. Zhang, Z. Li and J. Yan, Chin. Phys. L&t., 3 (1986) 9. G. G. Guilbault, Pructical Fluorescence, Marcel Dekker, New York, 1973, pp. 11-14. J. W. Bridges, in J. N. Miller (ed.), Stuncfurdr in FZuorescence Spectrometry, Chapman and Hall, New York, 1981, p. 75. C. A. Parker, Photoluminescence OfSoZutions,American Elsevier, New York, 1968, pp. 22&234. J. D. Winefordner, S. G. Schulman and T. C. O’Haver, in P. J. Elving and I. M. Kolthoff (eds.), Luminescence Spectrome@ in AnaZyticaZChemistry, Wiley, New York, 1972, pp. 81-103. J. N. Demas and G. A. Crosby, I. Phys. Chem., 75 (1971) 991. J. B. Birks, J. Res. NutI. Bur. Stand., Sect. A, 80 (1976) 389. J. B. F. Lloyd, in J. N. Miller (ed.), Standards in Fluorescence Spectromeby, Chapman and Hall, New York, 1981, p. 27. G. G. Guilbault, Practical FZuorescence, Marcel Dekker, New York, 1973, pp_ 17-19. I. L. Arbeloa, J. Chem. Sot., Faraday Trans. 2, 77 (1981) 1725. R. C. Weast (ed.), CRC Handbook of Chemisny and Physics, CRC Press, Boca Raton, 1979. J. W. Vaughn, in J. J. Lagowski (ed.), The Chemistry of Non-Aqueollr SoZvents, Vol. II, Academic Press, New York, 1967. R. S. Drago and K. F. Purcell, in T. C. Waddington (ed.), Non-Aqueous Solvent Systems, Academic Press, New York, 1965. G. J. Blanchard and M. J. Wirth, J. Phys. Chem., 90 (1986) 2521. J. Figueras, J. Am. Chem. Sot., 93 (1971) 3255. I. L. Arbeloa and K. K Rohatgi-Mukherjee, Chem. Phys. I&t., 128 (1986) 474. I. L. Arbeloa and K. K_ Rohatgi-Mukherjee, Chem. Phys. I&t., 129 (1986) 607. M. J. Kamlet and R. W. Taft, J. Am. Chem. Sot., 98 (1976) 377. R. W. Taft and M. J. Kamlet, J. Am. Chem. Sot., 98 (1976) 2886. M. J. Minch and S. S. Shah, J. Org. Chem., 44 (1979) 3252. M. D. Joesten, J. Chem. Educ., 59 (1982) 362. E. G. McRae, Spectrochim. Actu, I2 (1958) 192. N. Mataga and T. Kubota, Molecular Interactionsand Electronic Spectra, Marcel Delcker, New York, 1970, Chapters 6-8. K. Izutsu, Acid-Base Dissociation Constants in DipoZarAprotic Solvents, Blackwell Scientific Publications, Boston, 1990.
358 52 A. Streitwieser, Jr. and C. H. Heathcock, Introduction to Organic Chemishy, Macmillan, New York, 1976, Chapters 11 and 31. 53 D. W. Meek, in J. J. Lagowski (ed.), ir;he Chemisby of Non-Aqueous Solvents, Vol. I, Academic Press, New York, 1966, Chapter 1. 54 N. Nizamov, K. U. Umarov, R. Kh. Dzhumadinov and A. K. Atakhodzhaev, Opt. S’ctrosc., 54 (1983) 600. 55 G. S. Beddard, T. Doust, S. R. Meech and D. PhiIlips, J. Photo&em., Z7 (1981) 427. 56 N. S. Bayliss and E. G. McRae, J. Phys. Chem., 58 (1954) 1002. 57 A. A. Lamola, in P. A. Leermakers and A. Weissberger (eds.), Energy Transfer and Organic Photochemistry, Interscience, New York, 1969, Chapter 2. 58 S. G. Schulman, in E. L. Wehry ted.), M 0 dem Fluorescence Spectroscopy, Vol. 2, Plenum, New York, 1976, Chapter 6. 59 I. B. Berhnan, Handbook of Fluorescence Spectra of Aromatic Molecules, Academic Press, New York, 1971. 60 L. E. Cramer and K. G. Spears, J. Am. Chem Sot., ZOO (1978) 221. 61 G. C. Pimentel, J. Phys. Chem., 79 (1957) 3323. 62 P. B. Bisht, H. B. Tripathi and D. D. Pant, J. Photo&em Photobiol. A: Cbem., 58 (1991) 295. 63 H. Shizuka, Act. Chem. Res., 18 (1985) 141. 64 M. Vogel, W. Rettig, U. Fiedeldei and H. Baumgtirtel, Chem. Phys. Let& 148 (1988) 347. 65 R. Gvishi and R. Reisfeld, Chem. Whys. L&t., 156 (1989) 181. 66 E. M. Kosower and D. Huppert, Annu. Rev. Phys. Chem., 37 (1986) 127. 67 M. Vogel, W. Rettig, R. Sens and K. H. Drexhage, Chem. Phys. Letf_, Z47 (1988) 452. 68 M. J. Snare, F. E. Treloar, K. P. Ghiggino and P. J. Thistlethwaite, I. Photo&em., 18 (1982) 335. 69 R. S. Becker, Theory and InteTretation of Fluorescence and Phosphorescence, Wiley-Interscience, New York, 1969, Chapter 17. 70 T. Forster, Chem. P&s. Lett., 17 (1972) 309. 71 0. Valdes-Aguiler and D. C. Necker, Act. Chem. Res., 22 (1989) 171. 72 B. C. Burdett, in E. Wyn-Jones and J. Gormally (eds.), Aggregation I3ocesse.s in Solution, Elsevier, New York, 1983. 73 E. Coates, J. Sot. Dyers Colour., 85 (1969) 355. 74 W. G. Herkstroeter, P. A. Martic and S. Farid, J. Am. Chem. Sot., 212 (1990) 3583. 75 K. Bergmann and C. T. O’Konski, J. Phys. Gem., 67 (1963) 2169. 76 R. W. Chambers, T. Kajiwara and D. R. Keams, J. Phys. Chem., 78 (1974) 380. 77 J. E. Selwyn and J. I. Steinfeld, J. Phys. Chem., 76 (1972) 762. 78 M. Gratzel and J. K. Thomas, J. Am. Chem. Sot., 95 (1973) 6885. 79 K. A. Zachariasse, N. V. Phuc and B. Kozankiewicz, J. Phys. Chem., 85 (1981) 2676. 80 K. Y. Law, Photo&em. Photobiol., 33 (1981) 799. 81 P. Mukerjee and J. R. Cardinal, J. Phys. Gem., 82 (1978) 1620. 82 K. Kalyanasundaram and J. K Thomas, J. Am Chem. Sot., 99 (1977) 2039. 83 M. E. Diaz Garcia and A. Sanz-Medel, Talantu, 33 (1986) 255. 84 H. Chiang and A. Lukton, J. Phys. Chem., 79 (1975) 1935. 85 K. S. Birdi, H. N. Singh and S. U. Dalsager, I. P&x C&m., 83 (1979) 2733. 86 L. J. Cline Love, J. G. Habarta and J. G. Dorsey, Anal. Chem., 56 (1984) 1132A. 87 W. J. Dressick. B. L. Hauenstein, Jr., T. B. Gilbert, J. N. Demas and B. A. DeGraff, .J. Phys. Chem., 88 (1984) 3337. 88 F. Podo, A. Ray and G. Nemethy, J. Am. Chem. Sot., 95 (1973) 6165. 89 N. J. Turro and I. F. Pierola, J. Phys. Chem., 87 (1983) 2420. 90 F. E. Bailey, Jr. and J. V. Koleske, PoZy(ethylene oxide), Academic Press, New York, 1976. 91 H. Singh and W. L. Hinze, AnaZ. Lett., 15 (1982) 221. 92 C. Tanford, Y. Nozaki and M. F. Rhode, J. Phys. Chem., 81 (1977) 1555. 93 S. J. Formosinho and M. M. Miguel, J Chem Sot., Faraduy Tmns. Z, 81 (1985) 1891.
343
Cresyl violet chemistry and photophysics in various solvents and micelles Stefan
J. Isak and Edward
Depa~ent
M. Eyring
of Chemirtry, Univeni~ of Utuh, SaIt Lake City, UT 84112 (USA)
(Received August 28, 1991; accepted January 16, 1992)
Abstract The chemical and photophysical aqueous micellar and organic
behavior of cresyl violet perchlorate was studied in aqueous, solutions. In aqueous micellar solutions the fluorescence
quantum yield of cresyl violet increases in sodium dodecylsulfate (SDS), whereas a dramatic decrease is observed in Triton X-100 (TX). The chemical and photophysical behavior of cresyl violet is strongly dependent on hydrogen-bond, electrostatic and acid-base interactions in solution, as well as un its hydrophobicity.
1. Inh-oduction used as a laser dye. It is a cationic Cresyl violet (CV’) (Fig. 1) is commonly chromophore which absorbs red light and has a fairly high fluorescence quantum yield &. The highest & for CV+ is observed in alcohols. A value for & of CV+ in dilute methanolic solutions (less than 10e6 M) of approximately 0.70 has recently been determined [l]. CV+ has also been applied in organic chemical sensors [2], solar energy conversion and energy storage [3], as a photosensitizer in energy and electron transfer reactions [4] and as a potential photosensitizer in photodynamic therapy [5]. Thus a clearer understanding of the chemistry and photophysics of CV+ would be of great interest.
Fig. 1. The two major resonance structures of cresyl violet perchlorate.
lOlO-6030/92/$5.00
0 1992 - Elsevier Sequoia. All rights reserved
344 CV+ is an oxazine-type dye. This class generally has properties similar to the related xanthene and rhodamine dyes [6-S]. CV+ possesses an insignificant triplet population on excitation and is found to be quite photochemically stable [6, 91. CV+ is a highly rigid, planar molecule with the amine hydrogens lying in the plane of the molecule [6]. The amino groups are relatively acidic [6, 81, and CV+ and related dyes have been observed to be deprotonated in basic solution 15, 8, 10-K?]. The major sources of non-radiative decay arise from the small energy difference between the ground and excited states [6] and the high frequency N-H vibrations of the amino groups [B, 131. As the energy gap between the ground and excited states decreases, these vibrations become more important in non-radiative decay [8]. Because of its structure, CVf has the potential to dimerize and does so extensively in aqueous solution where dimer formation is assisted through hydrogen bonding with water [14, 151. The hydrophobic@ of CV+ is also a contributing factor to dimer formation [14, 161. Few studies concerning CV + in particular have appeared in the literature. More often than not, CV+ has been included as part of a more general study and, as a consequence, different conclusions regarding its behavior have been reached. Our initial interest in CV+ was to use it as a fluorescence standard, and thus its fluorescence quantum yield was characterized in methanolic and aqueous solutions as a function of concentration [l]. In considering CV+ for use as a laser dye, we would like to determine its behavior in aqueous micellar solutions, since the use of aqueous micellar solutions of laser dyes has a number of advantages over other solvents [6, 171. With these considerations in mind the concentration-dependent r#+ values of CV+ in aqueous micellar solutions of three common surfactants were investigated. At the same time the absorbance and fluorescence characteristics of CV’ in a variety of solvents were investigated to provide a more comprehensive understanding of the chemistry and photophysics of CV+.
2. Experimental details 2.1. Absolute fluorescence quantum yields of cresyl violet in methanol Absolute values of & for CVf in methanol were determined using photothermal beam deflection photoacoustic spectroscopy (PBDPAS) [l]. The theory of this technique has been treated by both Sullivan and Tam [l&l and Terazima and Azumi [19]. The experimental set-up is shown in Fig. 2. The pump laser was a Quanta Ray DCR-2 Nd:YAG laser from which the frequency-doubled h = 532 nm output was used to pump a Lambda-Physik FL2002 dye laser. The laser dye employed was rhodamine 640 perchlorate. An excitation wavelength of 610 nm from the dye laser was used as the pump beam. The pulse repetition rate was 10 Hz and the energy of the pulse was varied between 0.1 and 1.0 CLJ.A 3 mW diode laser (A=780 run) or a 5 mW HeNe laser was used to supply the probe beam. Excitation and probe energies were adjusted using neutral density filters. The probe beam signal was detected with a photomultiplier tube (Hamamatsu R928) wired for fast response [20]. The photoacoustic signals were averaged on a LeCroy 9400 oscilloscope and stored on disk by an IBM PC compatible computer. Sample cells were made of quartz with an optical path length of 1 cm. Measurements were made at an ambient temperature of 24.0&0.2 “C.
345
photomultiplier tube
probe
laser
m X.YBZ
platform
splitter
X.Y.2
platform
Fig. 2. Experimental arrangement for the thermal lens measurements. The output from an Nd:YAG-laser-pumped dye laser is directed through a diaphragm. An HeNe or diode laser probe beam is made collinear with the pump beam using a beam splitter. The positions of the pump and probe beams as well as their focusing can be additionally controlled with a lens. The pump and probe beams are then directed approximately 2.5 m beyond the cell using a series of prisms. The probe beam is then cut approximately in half with a razor edge and focused onto a photomultiplier tube (PMT) using another lens. A filter is placed in front of the PMT to filter out any light from the pump beam. The output from the PMT is displayed on an oscilloscope where the waveform may be plotted or stored on disk by a computer. (Reprinted from S. J. Isak, S. J. Komorowski, C. N. Merrow, P. E. Poston and E. M. Eyring, Appl. Spectrosc., 43 (1989) 419. Copyright 1989, Society for Applied Spectroscopy, Frederick, MD.) 2.2.
Absorbance and fluorescence measurements
A Perkin-Elmer MPF-66 fluorescence spectrophotometer was used for the fluorescence measurements. Emission spectra were corrected using a totally reflecting surface supplied by the manufacturer. Absorbance measurements were made at an ambient temperature of 24.0 I,!I0.2 “C using a Perkin-Elmer Lambda 9 spectrophotometer. 2.3. Refractive index measurements The refractive indices of the aqueous micellar solutions at ambient temperature were determined with a Bausch & L-omb Abbe-3L refractometer. The results are included in Table 1 (see Section 3). 2.4. Equations Equation (1) [21-281 was used to determine the absolute fluorescence quantum yield of CV+ in methanol from the PBDPAS measurements (1)
346
where hF is the mean fluorescence wavelength determined from corrected emission spectra. A, is the excitation wavelength, H,, is the non-radiative emission intensity of emission intensity of a non-luminescent the sample and H,* is the non-radiative reference. The non-luminescent reference sample used was CV* in KI-saturated methanol. KI completely quenches CV’ fluorescence and produces no changes in the absorbance spectra. The fluorescence quantum yields of CV+ in aqueous and aqueous micellar solutions were determined relative to the absolute & values of CV+ in methanol according to the following relationship [29, 301
where &,ref is the fluorescence quantum yield of a reference compound (in this case the absolute & of CV’ in methanol),A is the absorbance, nD is the index of refraction of the solvent and a is the area under the fluorescence peak. Equation (2) is valid as long as Beer’s law is obeyed and there is no inner filter effect [31-371 on the fluorescence. For CV+, eqn. (2) is only valid for concentrations below approximately lO-‘j M [l], Above this concentration of CV+, dimerization and/or inner filter effects cause problems. To compensate, solutions of matched absorbances were used to determine & values at CV+ concentrations above 10m6 M. 2.5. Maserials EU (Mallinckrodt) was recrystallized from water and subsequently dried and stored in a desiccator under vacuum. Rhodamine 640 perchlorate and cresyl violet perchlorate from Exciton, sodium dodecylsulfate and hexadecyltrimethylammonium bromide from Sigma, Triton X-100, N,N-dimethylformamide (high performance liquid chromatography (HPLC) grade), methylsulfoxide (HPLC grade), triethylene glycol dimethyl ether and tetraejhylene glycol from Aldrich, acetone (PHOTREX; spectrophotometric grade) and ethyl ether (anhydrous) from J. T. Baker and hexane (Omni-Solv; spectrophotometric grade) and methanol (Omni-Solv; spectrophotometric grade) from EM Science were all used as received. Ultrapure water was used from a Corning MP-190 water purification system connected to an MP-1 water still.
3. Results and discussion 3.1.
Solubilization
and
ground state characteristics of cresyl violet
Absorbance maxima, mean fluorescence wavelengths and fluorescence quantum yields for dilute aqueous, aqueous micellar and organic solutions of CV’ are listed in Table 1. A number of solvent electro-optical properties, as well as the apparent solubility of CV+ in these solvents, are also listed. The solvents are presented in the order in which the’absorbance maximum of CVc is observed to increase (i.e. shift to longer wavelengths). A brief glance at Table 1 suggests little, if any, correlation between the solubility or absorbance maxima of CVf and any electro-optical property of the solvents listed. However, closer inspection of the data yields a definite correlation between the solubility and absorbance maxima and a number of solute and solvent properties. These include the polar and electrostatic solute-solvent interaction, the hydrophobic@ of CV+, the
(CO.40)
617 620 621 624 628 626 627 629 629
589 (462) 593 (472) 594 598 599 601 603 607 607
‘Ref. 38. b20 “C. ‘Ref. 39. dRef. 40.
(0.40-0.67) (0.40-0.67) 0.67 0.52 (<0.40)
(=S94)
(=595) 623 624
480 (603) 482 585 589
0.22 (x0.40) ( < 0.40)
(<0.40) (<0.40) 0.40 0.40
(<0.40)
(5589)
473 (604)
Triethylene glycol dimethyl ether Ethyl ether N,N-Dimethylformamide Water Hexadecyltrimethylammonium bromide (10 mM) Acetonitrile Acetone Methanol Sodium dodecykulfate (70 mM) Tetraethylene glycol-water (50:50) Triethylene glycol dimethyl ether-water (50:50) Triton X-100 (30 mM) Tetraethylene glycol Methylsulfoxide Hexane
4F
Mean fluorescence wavelength (nn.0
Maximum absorbance wavelength (nm)
Solvent
High High High Insoluble
High High High High
Very low High ~10-~ M
Low (-~10-~ M)
Solubiliry
1.96d 0.294’
46.6d 1.89”
1.3360f 0.0004 1.458” 1.476” 3.96” 1.372” 0
1.342” 3.92” 1.357’ 2.88’ 1.326’ 1.70a 1.3355f 0.0006 0.34Y 36.2d 0.316” 20.7” 0.547” 32.6’
Dipole moment (debye)
1.352’ 1.15’ 1.427’ 3.82’ 1.333” 1.85’ 1.3336f 0.0005
$5 “C)
0.222’ 4.34a 0.796 36.7” 0.890” 78.5”
y2.S“C) ;25 “C) (cP)
Physical and spectroscopic properties of ditute (less than 10e6 M) solutions of cresyl violet in a number of solvents
TABLE 1
4
w
348
acid-base equilibrium of CV+ in solution and the hydrogen-bonding characteristics of the solvent. The importance of hydrogen-bonding interactions of the solvent with CV+ and related dyes has previously been recognized [41-44]. Of the pure solvents listed in Table 1, three can be considered to be amphiprotic, i.e. capable of acting as either hydrogen-bond donors (HBDs) or hydrogen-bond acceptors (HBAs) [45, 461. These are water, methanol (MeOH) and tetraethylene glycol (TEG). The remaining pure solvents, with the exception of hexane, may all be considered to be HBAs since none of them has a hydrogen available for proton-donating interactions. Hexane is nonhydrogen bonding (NHB). Since water is a component in all of the mixed solvent systems, the mixed solvents may also be considered to be amphiprotic. CV+ has a number of sites that may be involved in a variety of interactions. These are indicated in Fig. 3. The specific hydrogen-bonding interactions may be described as follows. The aromatic amine group may be involved in two types of interaction. The amine hydrogens can hydrogen bond with an HBA, e.g. with the oxygen atom in water or methylsufioxide (DMSO). The amine nitrogen may hydrogen bond with an HBD, e.g. with the hydrogen in water or MeOH. The heterocyclic nitrogen and oxygen may hydrogen bond with the hydrogen from an HBD. Finally, the hydrogen atoms of the protonated amine may interact with an HBA. Thus, in terms of hydrogen bonding, we might expect water, MeOH and TEG to provide the greatest solvating power. On the basis of ground state stabilization’we might expect the absorbance maxima of CV’ in these three solvents to be blue shifted with respect to the remaining solvents. This is because a red shift in the absorbance maximum with a decrease in hydrogen-bonding interaction is expected for CV+ 147-503. However, as is evident from Table 1, this is not the case. Also inconsistent in terms of hydrogen bonding is the low solubility of CV+ in water. Thus the solubility and stabilization of CV’ do not arise solely from hydrogen-bonding interactions of CV’ with the solvent. The limited solubility of CV” in water may be explained by the somewhat hydrophobic characteristics of CV+. Although the benz[a]anthracene skeleton of CV+ has been substituted by a number of polar and hydrophilic substituents, it still possesses considerable hydrocarbon character. The polar nature and electrostatic charge of C!V+ also play an important roIe in its solubilization and stabilization. Thus amphiprotic solvents provide extensive solubilization and stabilization for both CV+ and the perchlorate (ClO,-) anion. In the remaining solvents, which are all HBAs with the exception of hexane, extensive
H2y&; 8 H I
/
Ii
CIOF d
0
!
“\,/”
l-i2’---
.
/
0
\
i
“\,/”
H R
s
Hc;\ R
R
Fig. 3. Potential sites for hydrogen-bonding, electrostatic violet (see text for a discussion of these interactions).
and acid-base interactionsin cresyl
349
solubilization and stabilization of CV +, but not C104-, can be expected. The poorly solvated ClO.+- is thus potentially quite reactive. The degree of electrostatic interaction between solute and solvent should correlate fairly well with the dielectric constant or dipole moment of the solvent, the solvent with the largest value of the dielectric constant or dipole moment providing the most solubilizing and stabilizing power, Such electrostatic considerations may explain why the absorbance maximum of CV+ in acetonitrile (CE&CN) is blue shifted with respect to that in acetone, but do not explain why the absorbance maximum of CV+ in NJV-dimethylformamide (DMF) is significantly (greater than 100 nm) blue shifted with respect to that in CHsCN, which has a comparable dielectric constant, or to that in DMSO, which has a comparable dipole moment. The large blue shift of the absorbance maximum of CV4 in triethylene glycol dimethyl ether (triglyme), ethyl ether and DMF is owing to the deprotonation of CVc. Thus an acid-base equilibrium exists in solution between CV’ and the solvent. If the solvent is sufficiently basic and capable of accepting a proton, a transfer of one of the protonated amine hydrogens to the solvent may take place. Such a transfer would significantly stabilize CV+ and consequently result in a considerable blue shift of the absorbance maximum. The ethers and DMF are basic solvents [51-531 capable of accepting a proton. In the case of DMF, considerable resonance stabilization of the protonated form may be expected. Similarly, large blue shifts in the absorbance maximum, owing to the deprotonation of CVf, have been observed in a number of solvents, including DMF [6, 10, 12, 541. The addition of water or a small amount of acid to the deprotonated solutions of CVf restores the protonated form as illustrated by a red shift in the absorbance maximum back to approximately 600 nm. This can be clearly seen to occur with the triglyme and 50~50 water-triglyme solutions as indicated in Table 1. However, in at least one report [4], complex formation between CV+ and the basic solvent is suggested as opposed to the acid-base interpretation of the data discussed here. In some cases, the absorbance bands of both the protonated and deprotonated forms of cresyl violet can be observed. When this occurs one species is predominant. For the solvents in which this is observed the absorbance maximum of the minority form is indicated in parentheses next to the absorbance maximum of the predominant form. Some workers [13, 41,551 have attributed the spectroscopic characteristics of CV+ to interactions of the solvent solely with the heterocyclic nitrogen and/or oxygen in CV+. This seems unlikely based on the present discussion and other reports [6, 10, 12, 541. The solubility and absorbance data in Table 1 can thus be explained in terms of the hydrogen-bonding interactions between CV+ and the solvent, the somewhat interactions between CV+ and the hydrophobic characteristics of C!V+, electrostatic solvent and the acid-base equilibrium between CV+ and the solvent. While other contributions to the observed spectroscopic properties of CV+, such as the effect of the solvent refractive index on the absorbance maximum, are, no doubt, superimposed, they can be expected to be overshadowed by the contributions made by the acid-base, hydrogen-bonding and electrostatic properties summarized above [56]. The important solvent-solute interactions summarized above for CV+ facilitate an understanding of the difference in CV+ behavior in various solvents, even though it may be difficult to judge the relative importance of each in the overall solubilization and stabilization of CV+. For example, it should now be clear why CV+ is not soluble in hexane. First, hexane is non-polar with a very low dielectric constant. Thus little
350
Secondly, no hydrogen-bonding if any electrostatic interaction with CV + is expected. interactions are expected between CV+ and hexane, since hexane has no hydrogenbonding capacity, either as an HBA or HBD. Similar observations have been made for CV+ in other non-polar solvents [7, 415 In a comparison of the characteristics of CV+ in water, CH&!N, acetone and MeOH, the limited soIubility in water has already been attributed to the somewhat hydrophobic nature of CV+. The increase in the absorbance maximum of CV+ in the order water < CH&N < acetone < MeOH can be explained in terms of hydrogenbonding and electrostatic interactions between CVf and the solvent. Water and MeOH would provide maximal hydrogen-bonding interactions with CV+ serving as both HBAs and HBDs. The blue shift in water with respect to the other solvents is attributable to the strong hydrogen bonds formed and its high dielectric constant. The blue shift of CV’ in CI-&CN with respect to MeOH may be attributed to the fact that CH,CN is considerably more polar with a somewhat larger dielectric constant than MeOH. This results in a greater stabilization of CV+ in CH,CN than in MeOH. The absorbance maximum of CV+ in acetone is virtually the same as that in MeOH. While acetone is more polar than MeOH, it has a lower dielectric constant. The absorbance spectrum of CVf in acetone indicates the presence of both the protonated and deprotonated forms. As previously discussed, the deprotonation of CV+ results in considerable stabilization. These factors apparently contribute to the stabilization of CV’ to provide a ground state which is similar in energy (if not slightly lower than) in acetone to that in methanol. 3.2.
Excited state charactetitics
of cresyl violet
In all solvents, with the exception of DMF and the ethers, there is a significant overlap between the absorbance and fluorescence spectra of CV+. There is also a small Stokes’ shift. The largest Stokes’ shifts for CV’ are observed in water and aqueous hexadecyltrimethylammonium bromide (CTAB) solutions. As the absorbance and fluorescence spectra are progressively red shifted, a general decrease in the Stokes’ shift is observed, with the TEG and DMSO solutions exhibiting the smallest Stokes’ shifts. The range of the shift in the absorbance maximum is nearly twice that observed for the mean fluorescence wavelength hF. The absorbance spectra reveal a shoutder that is blue shifted with respect to the absorbance maximum. The fluorescence spectra are fairly good mirror images of the corresponding absorbance spectra. The fluorescence spectra were obtained by exciting the sample at 610 nm. The situation is quite different in DMF and the ethers. In these solvents deprotonated cresyl violet (CV) displays a very large Stokes’ shift and the absorbance and fluorescence spectra differ markedly. The absorbance spectra are centered around approximately 480 nm with fairly symmetrical peaks. The blue-shifted shoulders observed for CV+ are no longer present. The fluorescence spectra are quite broad, approximately 85 nm at full width at half-maximum (FWHM) compared with approximately 30 nm at FWHM for CV+. As a consequence, the fluorescence spectrum of CV overlaps the region of fluorescence of CV+ to a considerable degree. Owing to the lack of absorbance at 610 nm the fluorescence spectra of CV in these solvents were obtained by excitation at 490 nm. The fluorescence spectra display some structure that appears to be resolvable into four separate fluorescence peaks. Such an observation is unusual. The fact that excitation-wavelength-dependent emission spectra have been observed for similar deprotonated oxazine dyes [lo] suggests that the observed fluorescence spectra may depend on the excitation wavelength used. It is also curious that the observed fluorescence spectra are virtually the same in the three different solvents.
351
For the solvents in which b of C!V+ was determined, the highest value was observed in MeOH and the lowest in aqueous Triton X-100 (TX) solutions. The & values are listed in Table 1. Qualitative assessment of the J#+ values of C!V+ in acetone and CH3CN indicates that they are similar to that in MeOH and probably not less than the value in water. In the remaining solvents, the fluorescence is quenched with respect to that in water. The & values of CV in DMF and the ethers were more difficult to ascertain because of significant differences in absorbance. However, a & value less than that in water is expected [4, 10, 54, 57, 58). The estimated values of & for CV+ and CV are listed in parentheses in Table 1 together with the measured vaiues. The relatively small Stokes’ shift and large overlap of the absorbance and fluorescence spectra indicate small configurational changes between the ground and first excited states of CV+ [59]. This is confirmed by the good mirror symmetry of the absorbance and fluorescence spectra, indicating negligible reorganization effects and small geometry changes between the ground and first excited states [60]. The larger shifts observed for the absorbance maxima compared with I\Findicate that the ground state may be more polar and more solvent stabilized than the first excited state. This is reasonable if the large blue shift in absorbance on going from CV+ to CV is indicative of a general increase in the basicity of CV’ on excitation [SE&61-J. If this is the case, the importance of electrostatic and acid-base interactions of the solvent with the protonated amine group probably decreases on excitation. Such an interpretation also explains the larger shifts observed for the absorbance maxima compared with & and the general decrease in the Stokes’ shift as the absorbance and fluorescence spectra are progressively red shifted. For example, in a comparison of the data for water and DMSO, we would expect water to provide the most significant hydrogen-bond interactions and also significantly stronger electrostatic interactions because of the high dielectric constant. This results in a blue shift of the absorbance spectrum in water relative to that in DMSO. If, on excitation, some of the ground state interactions become less important, a larger energy difference is expected for water than for DMSO and is expressed in the larger Stokes’ shift. A similar observation in rhodamine 6G solutions has been explained as being due to orientational reorganization of nearby polar groups f17]. Such reasoning is also consistent with the conclusions reached regarding solvent effects on merocyanine dyes [49]. The absorbance spectrum of CV+ displays a blue-shifted shoulder. This shoulder disappears in the case where CV is formed. This blue-shifted shoulder has been ascribed to steric hindrance produced between the hydrogen of the amine group and the hydrogen of the phenyl group located nearby [63 or to two different oscillator groups within CV+ IlO] which are nearly degenerate. On deprotonation of CV+ the major resonant forms pictured in Fig. 1 will lose their resonance and the corresponding isomers will be formed. However, one of these isomers will still have the amine and phenyl groups cis to one another so that steric hindrance may still occur. There is no blue shoulder in the absorbance spectrum and the argument favoring similar oscillator groups in CV+ IS - more persuasive. As a brief aside, there appears to be some debate over whether the excitation and/or relaxation of CV+ may involve a charge transfer (CT) process [4, 58, 62, 651. A detailed discussion is beyond the scope of this report and may be found elsewhere [48,50,58, 661. However, it should be clear that the unprotonated amine or heterocyclic amine of CV+ may act as an electron donor and the protonated amine group may act as an electron acceptor. Such CT processes are expected to quench signscantly any fluorescence [50,58,62+4,66] and are not consistent with the significant fluorescence
352
observed. Thus it appears that a CT process is not operative. This lack of a CT process may be attributed to the greater rigidity of the aromatic amine groups of CV’ compared with the more bulky alkyl amino groups found in other oxazine and similar dyes for which CT processes have been observed 18, 64, 67, 681. CT processes do appear to be operative in the case of CV. These CT processes are a consequence of the acid-base equilibrium that can exist between CV+ and CV in solution. Two CT mechanisms have been proposed. The first mechanism involves the formation of a CT complex [4]_ CV is characterized by a significant blue shift in the absorbance maximum with respect to CV’ and its fluorescence is greatly diminished, displays a large Stokes’ shift and overlaps, to some extent, the region where CV+ is observed to fluoresce. In ref. 4, the large blue shift of the absorbance maximum of CV+ in amines is attributed to the formation of a CVf-amine complex. The large Stokes’ shift of the fluorescence spectrum of the complex and its coincidence with the fluorescence spectrum of uncomplexed CV+ are attributed to the similar fluorescence characteristics of uncomplexed and complexed CV+. Some of the complexed CV+ has been observed to dissociate on excitation. This dissociation is suggested to be owing to a smaller K, value in the excited state which results in a weakening of the CT interaction in the CV+-amine complex. Dissociation of the complex on excitation is indicative of a more basic excited state and thus of an increase in the pK, value of CVf on excitation as has previously been suggested. The other CT mechanism involves the transfer of a proton between CV+ and the solvent [lo, 11, 54, 651 and has already been discussed to some extent. In basic proton-accepting solvents a proton may be transferred from the protonated amine group of CV’ to the solvent. The stabilization obtained on formation of CV results in a considerable blue shift of the absorbance spectrum. Evidence suggesting that CV’ becomes more basic on excitation was discussed in the preceding paragraph. pK, changes between the ground and first excited states can be very large, with differences of 5-10 pK, units [SO, 58, 691 not uncommon. Thus, on excitation, it is possible for CV to become basic enough to be reprotonated to some extent, if not fully. Such a proton transfer process is expected to quench considerably any fluorescence [50, 58, 62-64, 661. We might also expect the observed fluorescence spectrum of CV to be similar to that of CV’. Both expectations are consistent with the data. The fluorescence is quenched in all cases for CV, there is considerable overlap of the fluorescence spectrum of CV with that of CV+, and the large Stokes’ shifts for CV indicate a ground state more stable than that of CV+, but a similar excited state. Proton transfer may also explain the fluorescence spectra observed for CV. The fluorescence spectra are broad and can be resolved into four fairly equivalent peaks. Two of these peaks are identical in intensity as are the other two. The differences in intensity are less than 10%. The slightly less intense peaks are blue shifted relative to the slightly more intense peaks. Thus it seems that there are two pairs of peaks. How they may be related is not clear. From Table 1 it can also be seen that the & value of the two slightly more intense peaks is always the same, while the hF value of the slightly less intense peaks varies with the solvent. There are two major resonance forms of CV+ (see Fig. 1) and, on deprotonation, this resonance is lost and the two corresponding isomers are formed. If steric hindrance is not a factor, a 50:50 mixture of the isomers may be expected. If the isomers of CV are fluorescent, as is CV+, it is possible to have four fluorescing species in solution at the same time. The observability and intensity of each fluorescent species are clearly dependent on the kinetics of all the excitation and de-excitation processes involved. The kinetic considerations involved have been considered elsewhere 150, 581.
353 The two isomers of CV are nearly, but not exactly, the same, just as the two different oscillator groups [lo] observable in the absorbance and fluorescence spectra of CV+ are nearly, but not exactly, the same (as discussed above). Such a difference is most clearly manifested in the excited state. The stabilizing power of the solvent may be similar for both the ground and excited states of CV. In this case we would expect to observe solvent-dependent IF values as is the case for the slightly lower intensity fluorescence peaks of CV. On the other hand, if CV is reprotonated on excitation, solvent interactions may be reduced and become less important as in the case of CV+. Such reasoning is supported by the generally solvent-independent AF values observed for CV+. On reprotonation of the CV isomers, resonance will be re-established. Although individual absorbance and fluorescence peaks of the resonance forms were not observed for CV+, it is apparent that some distinction is possible as illustrated by the shoulder observed in the absorbance and fluorescence spectra of CV+ . Such small energy differences may be manifested in the fluorescence spectrum of CV. The stabilization provided by DMF and the ethers may easily allow for a distinction between the two species. This suggestion is supported by a trend in the data indicating a decrease in the peak width (FWHM) of the fluorescence with an increase in the blue shift of the fluorescence spectrum, so that in an extremely stabilized environment, the two peaks may be resolved. Finally, the variation in 6 of CV + in different solvents merits discussion. While proton transfer appears to be predominantly responsible for the quenching of the fluorescence of CV, this is not the case for CV’. For CV+, hydrogen-bonding and electrostatic interactions can be expected to be important as well as any process that may contribute to an increase in the non-radiative decay rate. For example, the following factors all provide possible explanations for the greater quenching of CV + in DMSO than in MeOH. First, as a general rule, the rate of non-radiative decay increases with decreasing energy difference between the ground and excited states [68]. Secondly, and perhaps more likely, there is quenching of the fluorescence by the C104- counter-ion. From Table 1 it is apparent that little stabilization of CV+ is obtained through its interaction with DMSO. At the same time, owing to the nature of polar aprotic solvents, the C104counter-ion is poorly solvated and rendered potentially quite reactive. Thus the poor solubilization and stabilization will lead to a more tightly associated CV’-C104interaction in which the ClO,ion may quench the fluorescence of CV +. Finally, it is possible that the sulfur atom of DMSO may serve as an effective fluorescence quencher. The fluorescence quenching ability of sulfur, primarily through increased triplet yields, has been reported [5, 61. For example, the fluorescence of thiazines is quenched with respect to oxazines, because of the presence of a sulfur atom in the molecule. While most instances appear to involve intramolecular sulfur quenching, it is possible that an efficient interaction of the sulfonyl oxygen of DMSO with the cationic amine nitrogen will provide sufficient conjugation of DMSO with the r electronic system of CVc to result in effective quenching, where sulfur acts as the electron acceptor. The fluorescence of CV+ in water is also quenched with respect to that in MeOH for a variety of reasons. One contributing factor is the formation of non-fluorescent dimers in water as noted above. The importance of N-H vibrations in the non-radiative deactivation of CV+ has also been mentioned. Additional vibrational energy loss may occur for C!V+ through 0-Hvibrations with hydrogen-bonded water. The CV+-hydrogen bonds are expected to be both strong and efficient. Such additional energy losses have been suggested by Fiirster [70] and observed in the case of rhodamine B 1681.
354
3.3. Interaction of cresyl violet with aqueous micellar solutions The variation in & of CV* in water and aqueous micellar solutions of CTAB, sodium dodecylsulfate (SDS) and TX is presented in Fig. 4. The & values for the dilute micellar solutions of CV+ are listed in Table 1. From Fig. 4 it can be seen that the concentration-dependent & behavior of CV+ in CTAB is virtually identical to that in water. In SDS micellar solutions, an enhancement in the fluorescence of approximately 25% is observed, as opposed to a decrease in & of approximately 50% in micellar TX solutions. There is an increase in the observed absorbance maximum and & in the order water
concentration
Fig. 4. Concentration-dependent aqueous CXAB
[Id]
fluorescence
quantum yields of cresyl violet in aqueous (v),
(Cl), aqueous SDS (A) and aqueous TX
(0)
solutions.
355 the interaction between C!V+ and CI’AB which takes place in the ground state is diminished in the excited state. Interaction of CV+ with the Gouy-Chapman layer of the micelle may account for this. The Gouy-Chapman layer consists of a high density of counter-ions of the polar CX4B head groups, in other words Br- and C104-. On excitation, CV+ becomes more basic and interactions with such negatively charged species will diminish. It is interesting to note that no quenching of the fluorescence of CV* is observed as a consequence of its potential for interaction with Br-, a fairly effective quencher. Perhaps the interaction of CV+ with the Gouy-Chapman layer is inconsequential or CVc, Br- and C104are effectively shielded from one another as a consequence of the intense soIvating power of water. SDS is an alkyl sulfate anionic detergent. The negatively charged sulfate head group is counterbalanced by sodium Na +. There is no evidence for CV+ dimerization in the SDS micellar solutions. The polar micelle-water interface of SDS [7&82] and the complementary electrostatic charges of CV+ and SDS both favor a strong interaction between the two. In addition to electrostatic interactions, some hydrophobic interaction of the hydrocarbon-like portions of CV+ may be expected with the long hydrocarbon tails of SDS. Consequently, it is possible that CV+ may be localized at the micelle-water interface. In general, solute-micelle interactions become stronger as the solute becomes more hydrophobic [47, 83-861. The hydrophobic@ of CV+ has already been discussed. The hydrophobicity of CV+, the observed loss of dimerization of CV+ in SDS and the significant shift in the absorbance maximum of CV+ in SDS with respect to that in water all indicate a strong interaction of CV+ with SDS. Hydrogen-bonding interactions between CVf and water will be diminished on interaction of CV+ with the micelle. This is reflected by the red shift in the absorbance maximum of CV+ in SDS with respect to that in water, consistent with the observation that a decrease in hydrogenbonding interaction is expected to produce a red shift in the absorbance spectrum of CV+_ The lack of a significant difference in the & values between C!V+ in water and SDS may be explained as previously for CTAB. The & value of CV+ in SDS, while enhanced with respect to that in water, is still quenched with respect to the value in MeOH. So, although the fluorescence may be enhanced owing to the elimination of dimers and perhaps other factors, other quenching processes are still active or may be generated as a consequence of the interaction_ Two possibilities for quenching include strong hydrogen-bond interactions between CV* and water in the micellar solutions and energy loss via a strong electrostatic interaction between CV+ and the polar sulfate head group of SDS. TX is a non-ionic detergent and consists of a bulky alkyl phenyl head group and a long polyoxyethylene (POE) chain which terminates in a hydroxyl (-OH) group. This renders the POE chain amphiprotic in nature with the ether oxygens capable of serving as HBAs and the -OH group as an HBD or HBA. The micelles of TX differ from those of CT’AB and SDS. The long POE chains of TX are polar and thus extend out into the surrounding aqueous environment, whereas the long hydrocarbon taits of CTAE3 and SDS are located within the interior of the micelle. The long POE chains of TX have been found to be quite flexible and may interact with one another [87-901. Evidence for a strong interaction of CV+ with TX is given by the significant red shifts observed in both the absorbance maximum and & with respect to those obtained in water, the elimination of any evidence of dimerization of CV* and a decrease in the h value of CV+ with respect to that obtained in water. Reports of fluorescence quenching in micellar solutions are not common. In one such report [91], quenching was suggested to be attributable to impurities in the TX used. This possibility was
356
ruled out in our results when enhanced fluorescence was observed in other dye systems in our laboratory [9] using the same source of TX. On the basis of the observations and discussion of CV+ up to this point, an interaction of CVf with the POE chain seems quite likely. To test this hypothesis, the absorbance and fluorescence of CV+ were studied in triglyme (an HBA), TEG (amphiprotic) and SO:50 aqueous mixtures of each. The results are listed in Table 1 and a strong interaction, accompanied by a decrease in the & value of CV+, is indicated in each case. In triglyme, the ether oxygens are apparently basic enough to deprotonate CF. The addition of water, as in the 50~50 aqueous triglyme mixture, restores CV+ . However, compared with triglyme the behavior of CV+ in TX much more closely resembles that in TEG, indicating that interaction with the terminal -OH group is very important. The extent of interaction of CV’ with the ether oxygens of TEG is not clear. The observation that the values of the absorbance maximum and & of CV+ in TX lie between the values obtained in TEG and in 50:50 aqueous TBG solutions indicates the importance of water in the interactions. Clearly, a considerable amount of interaction between the POE chain and water is expected and has been observed experimentally 188, 921. Although the exact nature and extent of interaction of CV+ with the TX micelles are difficult to ascertain, it is clear that the POE chains are capable of providing hydrogen-bonding, electrostatic and acid-base interactions at least comparable with those of water in a less hydrophobic environment. As a consequence of the interaction of CV+ with the POE chains, its fluorescence is quenched, probably because of interaction with the ether oxygens. Additional channels of non-radiative decay are probably created by additional interactions between C!V+, the POE chains and water. The flexibility of the POE chains may also contribute to the efficiency of the nonradiative processes. These conclusions are consistent with strong POE chain interactions observed in other instances [87, 89, 931 and a report in which the observed decrease in the fluorescence of uranyl ions in TX micelles was attributed to an interaction with the POE chains [93].
Acknowledgments This research was funded by the Department of Energy, Office of Basic Energy Sciences. The experiments were initiated under a Biomedical Research Support Grant (USPHS No. RR07092). The authors thank J. D. Spikes for the use of his spectrophotometer and fluorometer.
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