J. Phys. Chem. 1988,92,45654568
Ftgure 18. A docription for the electron beam apparatus used for the electron beam exposures at atmospheric prc~rurerof dry nitrogen.
that fraction of the beam backscattered from the substrate. The charge density is defined as
Q = (1 + 1 ) q where q is the backscattering coefficient of the substrate. All of our experimental results are reported in terms of Q because it is a more precise measure of the number of electrons traversing a thin film. Q is readily converted to adsorbed dose using the methodology outlined previously." Elecrron Beam Exposure Appararus for Exposures a1 Almaipheric Pressures ofNilrogen. Electron beam exposure of the samples was achieved with a CR I SO Electron Processor (Energy
Sciences, Inc.. Woburn, MA) which allows exposures of the samples in an atmosphere of nitrogen. A diagram of the apparatus is shown in Figure 18. Basically, the instrument consists of a shielded conveyor that transports the sample under an clcctron beam. Electrons emanating from a rodlike filament are accelerated and subsequently exit through a Ti/AI alloy window to form a "planar" s h a d electron beam perpendicular to the direction in which the samples moves. The electron beam gun
4565
operates a t accelerating voltages between 150 and 175 kV. The dose delivered to the sample is controlled by adjusting either the beam current or the conveyor speed. The course of the electron beam induced decomposition was followed by using infrared spectroscopy. Infrared spectra of the solid thin film samples were coated on I-in.-diameter silicon substrates and recorded with a Perkin Elmer 580 IR spectrometer. The silicon substrates, obtained from Diode Corp., Framington, MA, are polished on both sides and are tapered (8-12 mil) to eliminate interference fringes. A bare silicon substrate was exposed to a 175-kV beam for doses up to 1000 Mrad with no new absorptions appearing. Dosimetry was performed using the aminopbenolmethane dye doped films produced by Far West Technology, Goleta, CA. The thickness of the dosimetric films was 2 mil and the absorbed dose was measured by recording the optical density a t 510 nm before and after exposure as reported earlier.'l A silicon wafer was placed beneath the dosimetric film to obtain the fraction of the absorbed dose from backscattering of the electron beam. The absorbed dose in the sample was obtained by using the equations for stopping power of the sample and d 0 ~ i m e t e r . l ~ Samples of irradiated material in sufficient quantities for product analysis were prepared by using a tray made from a metal block into which a well, 3 mil deep and of area equal to 160 cm2, was machined. A viscous solution of material was placed in the tray to flood the well and, subsequently, the sample thickness was brought to 3 mil by pulling a straightedge across the sides of the well and the remaining solvent allowed to dry. After the sample was exposed to the 175-kV electron beam it was collected and stored in capped bottles for analysis. Regis- No. 1,5610-94-6; 2, 114885-79-9; 7,879-15-2; 8, 108-39-4; 9, 71-43-2; 2-nitro-I-naphthol, 607-24-9; 2-amino-I-naphthol hydrochloride, 41172-23-0; sodium 2-diazo-l-naphth~quinone-5-sulfonate, 2657-00-3; 2-diazo-l-naphthaquinone-5-sulfonylchloride, 1 10928-59-1. (13) Pacansky,J.: Wang,C.; Waltman,R. J. 1.FImrimChem. 1986.32. 286. (14) Pacansky, I.: Waltman, R. I. 1.Rod. Curing 1986, 13. 24.
Emission Characteristics and Photostability of N,N'-Bis(2,5-di-tert-butylphenyl)-3,4:9,lO-perylenebis(dicarboximide) El-Zeiny M. Ebeid,* Samy A. El-Daly, Department of Chemistry, Faculty of Science, Tania Uniuersity, Tanta, Egypt
and Heinz Langbals Instiiut fur Organische Chemie der Uniuersitdt Miinchen, Karlstrasse 23, 0 - 8 0 0 0 Mtinchen 2, West Germany (Received: July 30, 1987: In Final Form: December 17. 1987)
The titled dye (1) shows very high fluorescence quantum yield values as well as photostability. The dye undergoes molecular aggregation both in the ground state (at a critical concentration of ca. 2 X IO4 mol dm-l) and in the excited slate (giving excimerlike emission at ca. 600 nm). It displays solvatochromism in both emission and UV-visible absorption spectra. The dye does not give laser emission upon pumping ethanolic solutions with a nitrogen laser (& = 337.1 nm, peak power of 100 kW) but acts as an efficient quencher of 1,4-bis(@pyridyl-2-vinyl)benzene(P2VB) laser dye. The quenching process obeys a static type mechanism. Equimolar mixtures of dye 1 and P2VB or 2,5-distyrylpyrazine (DSP) laser dyes also give no laser emission. With even higher peak power (200 kW, pulse duration of 800 LIS) a laser emission can be obtained from chloroform solutions of dye 1.
Introduetion The synthesis and spectral identification of several highly fluorescent and very highly stable perylene derivatives have been recently reprted.l Dyes of such unique characteristics are very
attractive in many areas, e.g., dye lasers, and solar energy conversion, and can serve as photosensitizers, photon counters, and ( I ) Langhals, H.Chem. Be,. 1985,118,4641.
0022-3654/88/2092-4565$01.50/00 1988 American Chemical Society
4566
The Journal of Physical Chemistry, Vol. 92, No. 15, 1988
Ebeid et al. TABLE I: Fluorescence Quantum Yield (4 ,) Values of Dye 1 in = 420 nm) Different Solvents (b,,,
solvent methanol ethanol
&f
solvent
0.99 0.96
cyclohexanol methyl iodide
&f
0.91 0.41
Laser activity has been tested by using a nitrogen laser (A, = 337.1 nm and peak power = 100 kW) as a pumping source.
3
Experimental Section The synthesis and purification of dye 1’ and P2VBZhave been described earlier. The sample of dye 1 used in this study contains the a and b atropic isomers in the ratio 70:30, respectively. Fluorescence and excitation spectra together with fluorescence quantum yields and photostability measurements were performed by using a Shimadzu RF 5 10 spectrofluorophotometer. Fluorescence quantum yields were measured relative to rhodamine 6 G as a reference standard for which +f = 0.96 (A,, = 420 nm in ethylene g l ~ c o l ) . ~Light intensity was measured by using ferrioxalate actinometry: and UV-visible absorption spectra were measured by using a Pye-Unicam SP 8000 spectrophotometer.
Results and Discussion The emission spectra (A,, = 500 nm) of dye 1 in chloroform are shown in Figure la. In a relatively dilute (ca. mol dm-3) solution, the dye shows a structured molecular emission that changes to an excimerlike emission in concentrated (ca. 0.03 mol dm-3) solutions. Molecular aggregation in dye 1 occurs also in the ground state at a critical concentration of ca. 2 X lo4 mol dm-3 as shown by electrical conductivity meas~rements.~Figure 1b shows the changes in the specific conductivity u (in S cm-’) as a function of concentration in ethanol a t 20 O C . The break in the curve corresponds to the ground-state critical concentration of molecular aggregation. The sharp break in the conductivity curve suggests that the aggregation is not just to dimers but rather to high oligomer aggregates. A concentration-absorption plot of dye 1 in ethanol gives a positive deviation from the Beer-Lambert law at an absorbance of 0.2 absorbance units and a concentration mol dm-3. of ca. 9 X The fluorescence quantum yields (4f,A,, = 420 nm) of unaggregated dilute solutions of dye 1 have been measured relative to rhodamine 6G after correction for the solvents’ refractive indexes.6 Table I summarizes the $Jf values, which are among the highest known. The observed low 4f value in methyl iodide reflects the external heavy-atom effect that causes radiationless deactivation of the first excited singlet state (SI)via intersystem crossing (isc). As expected from the very high 4f values, dye 1 is highly photostable. Photoirradiation (A,, = 365 nm and light intensity = 4 X einstein min-I) for ca. 60 min gave virtually no change in the mol dm-3) and concentrated emission spectra of both dilute ( ( mol dm-3) ethanolic solutions. The dye might possibly be used as a quantum counter having the versatility of high &values, photostability, and the insignificant changes in fluorescence intensities over the temperature range from 10 to 40 OC. Unfortunately, dye 1 solutions give no laser emission upon pumping with a nitrogen laser. The dye does not absorb significantly a t the pumping laser wavelength (A,, = 337.1 nm) and undergoes molecular aggregation at concentrations suitable for laser action (ca. mol dm-3). Molecular aggregation is accompanied by a substantial decrease in fluorescence efficiency, and the integrated areas of the fluorescence peaks decrease by ca. 94% as the concentration changes from to mol dm-3 in ethanol. Laser inactivity of the 2’,6’-xylidyl analogue has also been reported upon pumping with a nitrogen The extinction coefficient ( e ) values of dye 1 at 337 nm are relatively low, and in benzyl alcohol, for example, e = 4.6 X lo3dm3 mol-’ cm-’. This means that high material concentrations and pumping power are needed for optimum threshold. Laser emission from dye 1 solutions in chloroform has been observed upon pumping with a high-power (e200 kW) nitrogen laser of very short pulse duration (-800 p ~ ) No . ~ ~ more details are currently available. The laser activity of dye 1 in the presence of energy donors has also been investigated. Equimolar mixtures of dye 1 and P2VB or DSP give no laser emission upon pumping with N2 laser (peak power of 100 kW) despite the efficient energy transfer in these systems. The Stern-Volmer plots of P2VB quenching with dye 1 as a quencher at different medium viscosities and temperatures are shown in Figure 2. The second-order quenching rate constant (k,) increases
(2) Ebeid, E. M.; Sabry, M. M. F.; El-Daly, S. A. Laser Chem. 1985, 5, 223. ( 3 ) Harriman, A. J . Chem. SOC.,Faraday Trans. I 1980, 76, 1978. (4) Murov, S. L. Handbook of Photochemistry; Marcel Dekker: New York, 1973; pp 119-123.
( 5 ) Eicke, H. F. Top. Curr. Chem. 1980, 87. ( 6 ) Morris, J. V.; Mahaney, M. A.; Huber, J. R. J. Phys. Chem. 1976,80, 971. (7) (a) Sadrai, M.; Bird, G. R. Opt. Commun. 1984,52,62. (b) We thank Prof. Dr. N . Karl of Stuttgart University for carrying out this test.
a
= 2
. D
1
500
600
lo-‘
A (nm)
M/ 1 Figure 1. (a) Emission (&, = 500 nm) of dye 1 solutions in chloroform of concentration and (-) 0.03 mol dm-’. (b) Variations in specific conductivity u as a function of concentrations of dye 1 in ethanol at 20 OC. (-e-)
fluorogenic materials. In the present article, we report some more quantitative emission characteristics of the titled dye (dye 1) in relation to medium
b
a dye 1
effects. We also examine its laser activity both in solutions of its own and in the presence of energy donors, e.g., the recently reported2 laser dyes 1,4-bis(j3-pyridyl-2-vinyl)benzene (P2VB) and 2,5-distyrylpyrazine (DSP).
The Journal of Physical Chemistry, Vol. 92, No. 15, 1988 4567
Characteristics of a Perylene Excimer Dye
2
1 3
9
[a] x
15
I
320
300
280
21
400
7 (K)
MI/
I
\* %.-.I-.-
'.-
600
500
Figure 2. Stem-Volmer plots of P2VB quenching by dye 1 as a quencher (X, = 337 nm, X, = 420 nm). (a) Effect of viscosity: 0 , in 2-propanol;
+
0 1 mL of ethylene glycol 9 mL of 2-propanol; and 0, 3 mL of ethylene glycol 7 mL of 2-propanol. (b) Effect of temperature on the quenching rate constants.
+
600
400
400
A (nm) Figure 4. Changes in (a) absorption and (b) emission spectra of dye 1 in (-.-) ethanol, (X) ethanolic buffer of pH 1.7, and (---) 96% sulfuric
600
500
500
acid.
A (nm)
that could be of the charge-transfer (CT) character." The effect of medium on the electronic spectra of dye 1 has been studied. Dye 1 displays a solvatochromic effect in absorption, excitation, and emission spectra as shown in Figure 3. A linear correlation exists between the solvent polarity (AA and both the where Af is given as12J3 absorption and emission maxima A,
I
Af = ( E
600
500 A (nm)
400
300 600
500 A
400
300
(nm)
Figure 3. Effect of solvent (a) the absorption and (b) emission and diethyl ether (e = 4.335); ethanol (e = excitation spectra: (-e-) 24.30), (---) cyclohexane (e = 2.2); (X) benzyl alcohol (e = 13.1). e is the dielectric constant. Excitation spectra are taken by following emission at 540 nm. (.-e)
with increasing medium viscosity (Figure 2a) and decreasing temperature (Figure 2b). Given the solution lifetime of P2VB as 0.6 f 0.1 ns,* k in isopropyl alcohol a t 20 O C is calculated from the Stern-Volmer slope as 1.4 X lot3dm3 mol-' s-l. This value is nearly 3 orders of magnitude higher than the diffusion rate constant (kdin)in isopropyl alcohol given9 as 0.27 X 1O'O dm3 mol-' s-l. The viscosity and temperature effects together with the high kq values compared with km indicate a static quenching mechanism with subsequent ground-state complex formationt0
(8) Ebeid, E. M.;Kandil, S . H. J . Photochem. 1986, 32, 387. (9) Caivert, J. G.; Pitts, J. N., Jr. Photochemisfry; Wiiey: New York, 1966; p 627. (10) Penzer, G. R. An Introduction to Spectroscopy for Biochemists; Brown, S. B., Ed.; Academic: London, 1980; p 70.
- 1)/(2t
+ 1) - (n2 - 1)/(4n2 + 2)
where e and n are the dielectric constant and the refractive index of the solvent, respectively. The spectra are bathochromatically shifted as the solvent polarity increases, indicating an excited state that is more polar than the ground state." This effect together with the substantially high molar absorptivity (e = 95 000 dm3 mol-' cm-')' indicates a T-T* transition. The electronic spectra of dye 1 also show insignificant changes in both cationic (cetyltrimethylammonium chloride) and anionic (sodium dodecyl sulfate) micelles. It seems that the absence of charge separation and the bulky nature of the dye molecules prohibit its solubilization. The absorption spectrum of dye 1 is slightly bathochromically shifted (by ca. 5 nm) as the medium acidity increases (pH 1.7 or after flushing with HCl gas as shown in Figure 4a). The emission spectrum also decreases in intensity as shown in Figure 4b. It is concluded that protonation a t the heteroatoms is prohibited by steric factors. In highly concentrated sulfuric acid AnalaR, 96%), however, the dye changes from pink to deep blue with a substantial bathochromic shift of ca. 50 nm (Figure 4a), and the dye emission disappears. The original pink coloration, however, is reobtained on diluting the sulfuric acid. The blue coloration would be due to the formation of cationic species of dye 1. Structurally related tetraazaviolanthrone dyes give a similar bathochromic shift upon pr~tonation.'~ (1 1) Griffith, J. Colour and Constitution of Organic Molecules; Academic: London, 1976; p 76. (12) Brecht, E. Anal. Chem. 1986.58, 384. (13) Rao, C. N. R.; Singb, S.;Senthilnathan, V. P. Chem. Soc. Reu. 1976, 5, 297.
4568
J. Phys. Chem. 1988,92, 4568-4569
Acknowledgment. We thank Dr. B. D. Hockhart of the Queen’s University at Belfast and Dr. M. A. Salem of Tanta University of testing the ESR activity. We also thank Prof. Dr. N. Karl of the University of Stuttgart for the high-power nitrogen laser measurements.
This might indicate that the bathochromic shifts observed in the absorption spectra in dye 1 is due to simple protonation at the basic carbonyl centers. The blue species of dye 1 in concentrated H2SO4 give no electron spin resonance (ESR) signal, indicating the absence of radical species.
Registry No. 1, 83054-80-2.
(14) LukBE; Langhals, H. Chem. Ber. 1983, 116, 3524.
COMMENTS A fact that is so long that it appears to have been totally overlooked by S H R (and by their critics) is that many organic peroxides give significant yields of molecular hydrogen upon thermal decomposition in the absence of oxygen. The peroxides that thermalize to yield H2 have, without exception, a hydrogen atom attached to each peroxidic carbon atom; Le., they contain the HCOOCH moiety. In particular, H2has been shown to be a product from the thermal decomposition of di-sec-alkyl and of di-n-alkyl peroxides.21s22In many cases, the source of the H2 has been shown not to involve the intermediacy of H’ atoms. A concerted mechanism for H2 production which involves a cyclic six-membered ring transition state, originally proposed by Wieland and Wingler,’O has gained general acceptance.
Formation of Molecular Hydrogen in the Blmolecular Self-Reactlon of Hydroperoxyl Radicals In the Gas Phase’ Sir: Sahetchian, Heiss, and Rigny2 (SHR) have recently provided additional data which, so it is claimed, supports their earlier conclusion3 that molecular hydrogen is produced in ca. 8% of the bimolecular self-reactions of HOO’ in the gas phase at 150-200 OC and under atmospheric pressure; Le., k l / k 2= 0.086. In both HOO’
+
8%
/
HOO’
h
H2
+
H202
202
(1)
+
(2)
02
S H R generated HOO’ radicals by the thermal decomposition of di-n-alkyl peroxides, ROOR, in the presence of oxygen. The HOO’ radicals were presumed to be formed in the reaction RO’
+ O2
-
HOO’
+ product
A study of the thermal decomposition of di-sec-butyl peroxidelg has thrown some doubt on this concerted mechanism and even
(3)
where R = n-alk~l.~’It was found2 that the ratio [H2]/([H2] + [H202])was equal to 0.079 f 0.007 and was independent of the oxygen concentration. On this basis, it was correctly pointed out that atomic H’ could not be produced by the decomposition of ROOR or of RO’ (followed in either case by H’ + ROOR H2 product), nor could H2 be produced by the decomposition of RO‘. The earlier S H R study’ has been criticized on several grounds.e6 Thus, Glinski and Birks: using a different method for the gas-phase production of H O W , obtained k 1 / k 2< 0.0022 at 25 OC and 50 Torr of N1. These workers4perceptively suggested that it “would not be surprising if there were another source of hydrogen in this (the S H R ) complicated, radical system.” Furthermore, Baldwin et alS5have shown that the formation of H2 by reaction 1 would represent a chain termination in the slow reaction of H2 with O2at 500 OC and that this would be totally inconsistent with the results of their very complete modeling of this oxidation process and of the H2-sensitized decomposition of
-
+
(8) Formation of H2 by the thermal decomposition of a-hydroxy peroxides is very well doc~mented.”~ It was first detected in 1898 by Blank and Finkenbeiner9 in the reaction of H202and H2C0 in basic solution, the generation of H2 being traced to the intermediacy of bis(hydroxymethy1) peroxide.1° The formation of H2 during the thermal decomposition a-hydroxyalkyl peroxides occurs in the liquid phase,”’Jc” the vapor phase,l2*I3and even in the solid state.1° (9) Blank, 0.; Finkenbeiner, H. Ber. Dfsch. Chem. Ges. 1898, 31, 2979-2981. (10) Wieland, H.; Wingler, A. Jusfus Liebigs Ann. Chem. 1923, 431, 301-322. (1 1) Rieche, A.; Hitz, F.Ber. Dfsch. Chem. Ges. B 1929,62,2458-2474. (12) Style, D. W. G.; Summers, D. Trans. Faraday Soc. 1946, 42, 388-395. (13) Jenkins, A. D.; Style, D. W. G. J. Chem. SOC.1953, 2337-2340. (14) Wurster, C. F., Jr.; Durham, L. J.; Mosher, H. S. J. Am. Chem. Soc. 1958,80, 327-331. (15) Durham, L. J.; Wurster. C. F., Jr.: Mosher. H. S. J. Am. Chem. Soc. 1 9 d , 80, 332-337. (16) Durham, L. J.; Mosher, H. S. J. Am. Chem. SOC. 1960, 82, 4537-4542. (17) Durham, L. J.; Mosher, H. S. J . Am. Chem. SOC.1962, 84, 281 1-2814. (1 8) (a) The formation of H2 by the thermal decompositionof primar and secondary dialkyl peroxides in the liquid phaseI9” and in the vapor pha.d321-N has also been reported. (b) The yield of H2 from di-sec-butyl peroxide is very much less in the vapor phase than in s01ution.l~ (19) Hiatt, R.; Szilagyi, S. Can. J . Chem. 1969, 48, 615-627. (20) (a) Hiatt, R.; LeBlanc, D. J.; Thankachan, C. Can. J. Chem. 1974, 52, 4090-4094. (b) This paper records the highest yield of H2. Thermal decomposition of Ph2CHOOCHPh2at 11C-130 OC in solution gave approximately 90% H2 Ph2C0 by a nonradical mechanism. Photolytic decomposition in toluene yielded no H2. (21) Arden, E. A.; Phillips, L. J . Chem. SOC. 1964, 5118-5125. (22) Livermore, R. A.; Phillips, L. J . Chem. SOC.B 1966, 640-643. (23) Walker, R. F.; Phillips, L. J. Chem. SOC.A 1968, 2103-2106. (24) Thynne, J. C. J.; Yee Quee, M. J. J . Phys. Chem. 1968, 72, 2824-2831.
H202.
(1) Issued as NRCC No. 29203. (2) Sahetchian, K. A.; Heiss, A.; Rigny, R. J . Phys. Chem. 1987, 91, 2382-2386. (3) Sahetchian, K. A.; Heiss, A,; Rigny, R. Can. J . Chem. 1982, 60, 2896-2902. (4) Glinski, R. J.; Birks, J. W. J . Phys. Chem. 1985,89, 3449-3453; Ibid. 1986, 90, 342. (5) Baldwin, R. R.; Dean, C. E.; Honeyman, M. R.; Walker, R. W. J . Chem. Soc., Faraday Trans. I 1984,80, 3187-3194. (6) Golden’s suggestion7 that H2 was formed from H’ arising from the decomposition of alkoxy1 radicals, RO’ H‘ + product, can be ruled out by the SHR results reported in ref 2. (7) Golden, D. M., private communication quoted in ref 2.
+
-
0022-3654188 , ,12092-4568$01.50/0 0 1988 American Chemical Societv I
-
