naphthalene - VAPOR ABSORPTION SPECTRA AND OSCILLATOR...

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Unformatted text preview: VAPOR ABSORPTION SPECTRA AND OSCILLATOR STRENGTHS OF NAPHTHALENE, ANTHRACENE, AND PYRENE1 J. FERGUSON,2 L. W. REEVEs,3 AND W. G. SCHNEIDER ABSTRACT Vapor absorption spectra have been measured for naphthalene, anthracene, and pyrene using a modified Beckman DU Spectrometer. Vapor concentrations were determined using published vapor pressure data, and oscillator strengths were computed for a total of five tran— sitions in the three molecules. It was found that the vapor oscillator strengths are considerably lower than in solution by a factor greater than that predicted by classical theory. An absorp- tion region at about 41,000 cm.‘1 in anthracene is tentatively assigned as the 1B2u+ state calculated by Pariser to lie at 42,400 cn1._1. INTRODUCTION Although the electronic energy levels of aromatic molecules have been studied rather extensively both experimentally and theoretically, data relating to absorption intensities in the vapor are lacking. Nearly all measurements of absorption intensity have been made in solution and it is assumed that in every case the solvent modifies this intensity only slightly so that theoretical intensities can be compared with solution intensities directly, at least for allowed transitions. This assumption is derived from the classical theory of Chako (5) and is substantiated by the work of Almasy and Laemmel (1) on diphenyl. However, it was felt that this assumption should be investigated by a measurement of the vapor absorption spectra of a number of aromatic compounds of interest. We have therefore investigated the absorption spectra of naphthalene, anthracene, and pyrene and used published vapor pressure data to determine the vapor concentrations; finally calculating the oscillator strengths of a total of five electronic transitions in these mole— cules. Information about the intensity of distribution throughout each band system which is of theoretical interest (2) was also obtained. EXPERIMENTAL Research grade anthracene, naphthalene, and pyrene were further purified by chroma— tography using the method of Sangster (14). The spectra were taken in a silica cell, length 20 cm., containing a small quantity (a few milligrams) of the purified crystals. The cell was first pumped out to a residual pressure of 10—6 mm. Hg, then a small amount of argon was admitted (approximately % cm. Hg) before the sidearm of the cell was sealed off. This added argon enabled thermal equilibrium to be achieved more quickly. The sealed cell was mounted rigidly in a small air thermostat fitted with appropriate silica windows. A heating element covering most of the floor of the thermostat was used and ,,with the aid of natural convection within the lagged space a temperature variation of 3:03" C. could be achieved by using a contact thermometer and relay box. This temperature was measured in contact with the cell by means of a mercury in glass thermometer. No condensation on the windows was detected when the cell was cooled from the highest to the lowest temperature used in any experiment. Spectra were taken after 20 minutes at the chosen temperature. The spectra of anthracene were reproduced with different loadings of the cell. 1Manuscript reeeived May .9, 1957. Cantribntian fram the Divisian 0f Pare Chemistry, Natianal Research Canneil, Ottawa, Canada. Issued as N.R.C. Na. 4458. 2N.R. C. Pastdactorate Fella'w, new at Chemistry Department, University 0f British Calumbia, Vancouver, B.C. 3N.R.C. Postdoctarate Fellaw, new at 111 ellan Institute, University of Pittsburgh, Pa. 1117 1118 CANADIAN JOURNAL OF CHEMISTRY. VOL. 35, 1957 The absorption spectra were measured with a modified Beckman DU Spectrophoto« meter made semiautomatic by use of a constant speed motor, a Photovolt Photometer No. 520M, and a Varian G10 recorder, details of which will be published at a later date. Recorder tracings of the light transmission at each temperature were compared with tracings of the blank cell in order to compute optical densities. For anthracene and pyrene these blank runs were taken with the cell at room temperature where the. vapor pressures are too low to record absorption by these compounds. The blank determinations for the naphthalene spectrum were obtained by carefully removing, cleaning, and replacing the cell without changing the optical arrangement. The stability of the light source, a xenon high pressure a-c. arc (Osram XBO—l62), was checked by taking a number of blank runs before and after the absorption spectra were recorded. RESULTS Fig. 1 shows the vapor absorption spectrum of naphthalene in the region of the second transition. The extinction coefficients have been computed with the vapor pressure data of Bradley and Cleasby (4) (extrapolated) and Sears and Hopke (15), for the temperature 313° K., where the two sets of data agree. The oscillator strength was computed from the well-known relation f = 4.31x10—9fe dv. Whenf is oscillator strength, 6 is molar extinction coefficient, and 17 is frequency in cmfl. Table IA contains this quantity computed from two sets of vapor pressure data. Fig. 2 shows the vapor absorption spectrum of anthracene in the region of the first transition where extinction coefficients have been computed from the vapor pressure data of Sears and Hopke, and in Table IB the oscillator strength of this transition is calculated from three sets of vapor pressure data. Similarly Fig. 3 shows the extinction coefficient of the second transition in this compound from the vapor data of Sears and Hopke; the variation of the oscillator strength with the same vapor pressure data used above is shown in Table IC. Fig. 4 shows the vapor absorption spectrum of pyrene in the region of the second transition (pyrene has a weaker absorption region lying at slightly longer wavelengths) and Table ID contains the oscillator strength of this transition computed from the data of Bradley and Cleasby. Fig. 5 shows most of the absorption spectrum of the third transi- tion and Table IE contains the oscillator strength computed up to the cutoff in Fig. 5 from the vapor pressure data of Bradley and Cleasby. Table II summarizes the values of the oscillator strengths in the vapor and also com- pares them with solution values computed from absorption spectra kindly provided by Dr. R. N. Jones.* DISCUSSION The computation of the oscillator strength in each of the absorption transitions depends critically on the published vapor pressure data. The various studies of the vapor pressure of naphthalene show marked discrepancies in the range of our measurements (25—40° C.). The measurements of Sears and Hopke (14) made in our range agree (within 6%) with those of two other determinations (l7, 3). In spite of the fact that there seems to be some curvature in the Antoine equation near 20° C. in all of these results, we feel that they have the advantage of applying directly to our range of temperatures. The extrapolated results of Bradley and Cleasby (4) are lower, and while they might be less suspect from *J. Am. Chem. Soc. 67, 21.97 (1945); Can. J. Chem. 34, 1017 (1.956). We are indebted to Dr. Jonesfor leaning us the originals. MOLAR EXTINCTION COEFFICIENT X IO—3 MOLAR EXTINCTION COEFFICIENTX If)"3 MOLAR EXTINCTION COEFFICIENTXIO'4 FERGUSON ET AL.: VAPOR ABSORPTION SPECTRA WAVELENGTH I3) 2500 2600 2700 2800 04 l m I —7' _I__%I__I—I__I— 4| 4o 39 38 37 36 35 3000 3I00 3200 3300 3400 3500 3600 3700 in 0.5 2200 2300 2400 2500 45 44 43 42 41 4o WAVE NUMBERx 10-3 (cm-') FIG. 1. Naphthalene absorption spectrum, second transition. FIG. 2. Anthracene absorption spectrum, first transition. FIG. 3. Anthracene absorption spectrum, second transition. 1119 1 i i l l 1120 CANADIAN JOURNAL OF CHEMISTRY. VOL. 35, 1957 TABLE I TEMPERATURE, CONCENTRATION, AND OSCILLATOR STRENGTH Source of vapor Temperature, Concentration, Oscillator Average pressure data T° K. moles/liter strength,f f A. Second naphthalene transition Bradley and Cleasby 298.7 5.47*X10_° 0.063 (6.7°—20.7° C.) 302.2 7.51* 0.068 0.062 313.2 22.2* 0.055 Sears and Hopke 298.7 7.26X10‘G 0.047 (19°—35° C.) 302.2 11.5 0.044 0.049 313.2 22.2* 0.055 B. First anthracene transition Bradley and Cleasby 422.7 10.27*X10—5 0.012 (65°—80° C.) 414.7 5.74* 0.014 0.013 409.7 396* 0.014 402.2 2.22* 0.013 Sears and Hopke 422.7 7.77*X10‘5 0.016 (105°—125° C.) 414.7 4.43* 0.018 0.017 409.7 3.09* 0.018 402.2 1.77* 0.016 Nita, Selci, and Mornatau 422.7 5.31*X10‘5 0.023 (72°—100° C.) 414.7 3.09* 0.026 , 0.024 409.7 2.18* 0.025 402.2 1.26* 0.023 C. Second anthracene transition Bradley and Cleasby 364.2 8.38*X10_7 , 0.74 (65°—80° C.) 358.7 494* 0.86 0.80 Sears and Hopke 364.2 7.75" X 10‘7 0.80 0.85 (105°—-125° C.) 358.7 463* 0.91 Nita, Seki, and Mornatau 364.2 5.86X10“7 1.10 (72°—100° C.) 358.7 3.56 1.20 1.15 D. Second pyrene transition Bradley and Cleasby 390.9 446* X 10—6 0.107 (72°—85° C.) 384.7 2.75* 0.101 0.102 379.7 1.88* 0.098 E. Third pyrene transition Bradley and Cleasby 390.9 4.46* X 10‘6 0.1181“ (72°—85° C.) 384.7 2.75* 0.1281‘ 0.117 379.7 1.88* 0.1041‘ *Denotes extrapolation of vapor pressure data. TTo cut—of. the point of view of impurity, they were made in a region where measurements are more difficult. Two determinations of anthracene vapor pressure are in fairly good agreement (4, 15) but a third set (10) of results indicates a lower vapor pressure. Both the effusion method of Bradley and Cleasby (4) and the modified Rodebush gauge used by Sears and Hopke (15) appear to be reliable and the results agree with the less accurate determination by Stevens (16) using the intensity of the fluorescence as a measure of the concentration. There seems no reason why the effusion method used by Nita, Seki, and Mornatau (10) should be less accurate so concentrations calculated from these data were also used to compute the oscillator strengths. Pyrene vapor pressures were measured by Bradley and Cleasby (4), and there has been an extrapolation of their results for the purpose Of the present work. The concentrations computed for pyrene are therefore less reliable than for the Others but still indicate considerable intensification of the absorption intensity from vapor to solution. FERGUSON ET AL; VAPOR ABSORPTION SPECTRA 1121 WAVE LENGTH (11°) 2900 3000 3I00 3200 25 4 ' {g WAVELENGTH (11°) T 20- 25&—26‘29___2_7'L O ; 5 ,_ 5 ® 9 & 15r- l5 «J O D Z 9 p— 3 ; 10— lo E n: 5 O 2 5- 5 l l I I —l _| I 35 34 33 32 31 40 39 38 37 WAVE NUMBER x 10" (cm") WAVE NUMBER x10"(cm-‘) FIG. 4. Pyrene absorption spectrum, second transition. FIG. 5. Pyrene absorption spectrum, third transition. TABLE 11 COMPARISON OF OSCILLATOR STRENGTHS IN VAPOR AND SOLUTION Vapor Solution Compound Range of f Average f f Naphthalene 2nd transition 0049—0062 0.055 0.113 Anthracene lst transition 0013—0024 0.018 0.11:, 2nd transition 0.80 —1.15 0.93 1.56 Pyrene 2nd transition 0098—0107 0.10 0.300 3rd transition 0104—0128 0.12 0.265 (to cutoff) The question “Does ultraviolet absorption intensity increase in solution?” was asked and answered by Jacobs and Platt (8) for isoprene, cis-, and trans-piperylenes. They con- cluded that the correction factor is close to unity, a result previously obtained by Pickett and co-workers (12) for cyclopentadiene and cyclohexadiene. Almasy and Laemmel (1) "have extensively investigated the absorption spectrum of diphenyl using a photographic technique and have found a correction factor of 1.32, and more recently Price and Hammond (13) have found both an increase and a decrease over solution intensities for two transitions of benzene in the vacuum ultraviolet. As can be seen from Table 11 our results for naphthalene, anthracene, and pyrene show a marked intensifiCation of absorp— tion intensity in solution. Although there are variations in the vapor pressures measured by different authors these are relatively small and cannot account consistently, no matter how they are arranged, for the large changes observed between the absorption intensity of the vapor and the solution of these substances. Although there is a lack of published data concerning solvent intensification in aromatic hydrocarbon spectra some interesting results reported by Herrington and Kynaston (7) 1122 CANADIAN JOURNAL OF CHEMISTRY. VOL. 35, 1957 have established that the effect of the solvent is not simple and varies from molecule to molecule and also depends on the electronic transition in the molecule. These workers have reported the difference in molar extinction coefficient between cyclohexane and alcoholic solutions and between 1 N LiCl and alcoholic solutions for a number of aromatic molecules. Of most interest to the present discussion is the difference in the behavior of the 2600 A and 3800 A region of the anthracene spectrum. The former shows a marked intensification in going from alcohol to cyclohexane while the latter does not change much; both systems, however, are more intense in 1 N LiCl than in alcohol. As can be seen from Table II we also find a considerable difference between the 2600 A and 3800 A systems of anthracene, but in this case the solvent intensification is much greater for the 3800 A system. In view of the results of Herrington and Kynaston and those quoted in this paper there appears to be an obvious need for an investigation of solvent induced intensity changes. With intensity changes as large as those recorded above, the transition from unperturbed gas molecule to solvent perturbed molecule could be followed with the aid of pressure techniques similar to those developed by RObin and Vodar (18). These authors have noted an increase in the absorption coefficient of phenanthrene dissolved in compressed N2 and H2, but have worked at very high pressures; intensity changes should become apparent at much lower pressures. It is interesting to make a comparison between theory and experiment. Theory has aimed at increasing the accuracy of calculating transition energies and the recent work of Pariser (11) is evidence of this. However, calculated intensities differ considerably from experimental intensities and Table III summarizes the more recent theoretical values TABLE III THEORETICAL OSCILLATOR STRENGTHS or THEORETICAL f Molecule Pariser (ll) Moffitt (9) Ham and Ruedenberg (6) Naphthalene [I 0.256 0.38 0.11 0.11* Anthracene I 0.386 0.37 0.27 0.11* Anthraceue ll 3.23 4.8 3.09 145* Pyrene [I 0.012 Pyrene III 1.53 “Normalized" with raped to benzene. off. As can be seen from a comparison between Tables III and Table II, agreement is poor, particularly for pyrene. In connection with the theoretical work of Pariser (11) it is interesting to examine more closely the vapor absorption spectrum of anthracene (Fig. 3) in the region corresponding to the 2600 A system in solution. It is unlikely, in view of the shape of the absorption band in solution, that the low intensity region at about 41,000 cm.‘l belongs to the more intense part of the absorption and more likely , corresponds to a separate absorption transition. Pariser has calculated that a ‘32qu state should lie 2000 cm.‘1 below the 1133;“ state calculated at 44,400 cm.‘1 and it is possible that the absorption in the region around 41,000 cm.“1 corresponds to the former. This would explain the marked difference in the shape of the absorption curve in solution because both transitions would probably react differently tO solvent perturbation. Four of the five transitions studied have very similar red shifts in going from vapor to FERGUSON ET AL.: VAPOR ABSORPTION SPECTRA 1123 solution, amounting to about 1000 cmfl. The intense 2600 A system of anthracene, how— ever, shifts tO low energies by as much as 3000 cm.‘1 depending on the solvent. REFERENCES . ALMASY, F. and LAEMMEL, H. Helv. Chim. Acta, 33, 2092 (1950). ALTMANN, S. L. Proc. Phys. Soc., A, 63, 1234 (1950). BARKER, J. T. Z. physik. Chem. 71, 235 (1910). BRADLEY, R. S. and CLEASBY, '1‘. G. J. Chem. Soc. 1690 (1953). CHAKO, N. Q. J. Chem. Phys. 2, 644 (1934). HAM, N. S. and RUEDENBERG, K. J. Chem. Phys. 25, 13 (1956). HERRINGTON, E. F. G. and KYNASTON, W. J. Chem Soc. 3137, 3143 (1952). JACOBS, L. E. and PLATT, J. R. J. Chem. Phys. 16, 1137 (1948). 9. Momm, W. J. Chem. Phys. 22, 320 (1954). 10. NITA, J., SEKI, S., and MORNATAU, M. J. Chem. SOC. Japan, 71, 430 (1950). 11. PARISER, R. J. Chem. Phys. 24, 250 (1956). 12. PICKETT, L. W., PADDOCK, E., and SACKTER, E. J. Am. Chem. Soc. 63, 1073 (1941). PICKET’I‘, L. W. and HENRI, V. J. Chem. Phys. 7, 439 (1939). 13. PRICE, W. C. and HAMMOND, V. J. Trans. Faraday Soc. 51, 605 (1955). 14. SANGSTER, R. C. and IRVINE, J. W. J. Chem. Phys. 24, 670 (1956). 15. SEARS, G. W. and HOPKE, E. R. J. Am. Chem. SOC. 71, 1633 (1949). 16. STEVENS, B. J. Chem. Soc. 2943 (1953). 17. SWAN, T. H. and MACK, E. J. Am. Chem. Soc. 47, 2112 (1925). 18. VODAR, B. and ROBIN, S. J. Chem. Phys. 16, 996 (1948); 18, 1413 (1950). 9°“???WFQH i l i E l l l HEATS OF ACTIVATION IN ELECTRODE PROCESSES THE ELECTROCHEMICAL DESORPTION MECHANISM OF THE DISCHARGE OF HYDROXONIUM IONSl B. E. CONWAY AND J. O’M. BOCKRIS ABSTRACT Quasi—quantitative potential energy profile diagrams are estimated for the electrochemical desorption reaction Plnq_++Hnds,—l—eo' —> Hg of the electrolytic hydrogen evolution reaction at Ni, Cu, W, and Hg. The calculated heats of activation for this reaction increase with increasing heat of adsorption of H on the metal, the converse of the situation found for the simple proton discharge reaction H,,q'++eo_ —> Hnds.. The direction of the trend of calculated AHI values for electrochemical desorption, with heat of adsorption of H, is in agreement with the observed dependence of exchange currents on heat of adsorption for a series of transition metals. For systems where stoichiometric number data cannot be obtained, the calculations allow distinction between rate-determining proton transfer and electrochemical desorption reactions. The experimental groups of values for the H/D separation factor are interpreted. I. INTRODUCTION Three rate determining steps in the hydrogen evolution reaction (h.e.r.) from aqueous acid solutions are (4, 5, 6): H30++e0——> Hnds.(A)i H3O++Hnda,+eo“—> H2 (B)*; and Had5.+Had5, —> H2 (C). Criteria for a given rate determining step have been proposed in terms of Tafel slopes and stoichiometric numbers (5), but A and B can give the same b values (2.303(2RT/F)), although their stoichiometric numbers (v) differ. Sometimes v is not determinable, so that a further criterion must be available. Conway and Bockris have pointed out (9, 10) that the direction of dependence of i0 upon heat of adsorption of H at various metals should be different for reactions A and B since the heat of adsorption of H will appear in the expression for the potential energy of the state after the transition state in reaction A, but in the potential energy of the state before the transition state in B. The consequences of this concept are evaluated here quasi-quantitatively for reaction B. Previously it has been assumed, without proof (19), that the rates of reactions A and B will be affected in the same way by changes of heats of adsorption of H. This is shown to be theoretically invalid and in disagreement with experiment. II. POTENTIAL ENERGY DIAGRAM CALCULATIONS Inherent in potential energy surface calculations by the Eyring—Polanyi semiempirical method (17) is the difficulty of assigning the correct proportion of coulombic and dis— persion interaction energy. For electrode processes involving transitions between ions and molecules the usual assumptions of the contribution of coulombic energy are likely to be inapplicable. However, for reaction A in the her. and some others (18), e.g. that between alkali metals and alkyl halides (16, 21), use of a cross section of the potential energy surface (the potential energy profile) assuming a linear reaction co-ordinate gives reasonable agreement with experiment (32, 11). The configuration in the model to be considered (111a) is one in which...
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