NOTES sealed off a t the constriction while pumping. After placing the system in the oven described above, the absorption spectrum of the sodium-ammonia mixture was scanned repeatedly from 1000 to 350 mp in the temperature range from 188 to 205". For pure sodium or sodium-ammonia mixtures containing smaller initial concentrations of ammonia than 0.11 M , no absorption band was observed over the temperature range 188-205". However, for sodiumammonia mixtures in which the initial concentration of ammonia was 0.11 M , assuming the ideal gas law (the cell volume was 50 ml), a narrow symmetrical absorption band with no fine structure was observed a t 587 f 2 mp in two of three runs. A similar absorption band was observed in one of two runs where the initial ammonia concentration was 0.13 M . In each case where an absorption band was observed, it extended from 590 to 585 mp, aind the band width seemed to be independent of the temperature in the range from 188 to 205". During; a run, the intensity of the peak rose to a maximum, remained constant for about 15 min, and then decreased to zero once again. Where the initial ammonia concentration was 0.11 M , the maximum per cent absorptions (the transmittance multiplied by 100) for the three runs were 4,8, and O%, respectively. For an initial ammonia concentration of 0.13 M , the maximum per cent absorptions for two runs were 4 and O%, respectively. After each of the runs had been completed, the initial sodium film had disappeared completely. Prior to a run, these films appeared stable to the eye a t room temperature in the presence of ammonia. In contrast, potassium films were seen to disappear a t room temperature in the presence of ammonia gas at several atmospheres of pressure. The instability of the intensity of the band is not too surprising, in view of the well-known heterogeneous reaction between sodium and ammonia to form the amide. The constant absorption intensity which holds for approximately 15 min is probably due t o the establishment of a steady-state concentration of absorbing species in the gas phase. The rate of formation would be expected to depend on the quantity of sodium and ammonia in the cell and the temperature. The rate of decomposition would depend on the state of the Pyrex surface where decomposition takes place,s diffusion from the gas phase to the catalytic sites on the glass surface, and the temperature. The nonreproducibility of the Pyrex surface5 from run to run may account for the range of absorption intensities (from 0 to 8%) in the five different runs. The narrow absorption maximum at approximately 587 mp, which coincides with the 2Sl/2+ 2Pl/2, transition of the gaseous sodium atom, suggests that the absorbing species in the gaseous sodium-ammonia mixtures is one in which the 3s electron of the sodium atom remains essentially unperturbed. The absence of absorption in the 1000-mp region, which is characteris-
3335 tic of the solvated electron in liquid ammonia,1*2 indicated no appreciable charge transfer of the type reported in the liquid phase. At higher ammonia concentrations, where the gaseous solvent may be more liquidlike in structure, perhaps close to the critical point of ammonia, the possibility of observing chargetransfer processes may be more likely. Investigation of this effect at higher pressures could not be carried out in the present apparatus due to strength limitations. ( 5 ) I. Warshawsky,
J. CataZ., 3 , 291 (1964).
Kinetics of the Thermal a! -t ,f3 Polymorphic Conversion in Metal-Free Phthalocyanine
by James H. Sharp and Roger L. Miller Xeroz Research Laboratories, Rochester, New York (Received March 91, 1968)
The phthalocyanines are an important class of organic compounds which have been extensively studied for their pigment properties. Moreover, within the last decade these compounds have undergone intensive investigations with respect to their optical, magnetic, and electronic conduction properties. Because of their high thermal stability, the phthalocyanines offer many advantages in the field of molecular electronics. I n particular, since several of the phthalocyanines also exhibit photoconductive behavior, there has been wide interest in the use of these materials as photoconductors. These compounds are known to exist in several polymorphic forms. The p form is the most stable polymorph, and its detailed crystal structure has been reported by Robertson.2 Other polymorphs designated as the ( Y , ~ - ~ O y,l0-l1and x12 forms have been characterized by X-ray powder diffraction pattern and infrared or visible spectroscopy. Assourlo and Sidorov (1) See, for example, F. H. Moser and A. L. Thomas, "Phthalocyanine Compounds," Reinhold Publishing Corp., New York, N. Y., 1963. (2) J. M.Robertson, J. Chem. Soc., 615 (1935); 1195 (1936); 219 (1937). (3) G.Susich, Anal. Chem., 2 2 , 425 (1950). (4)F. R. Tarantino, D. H. Stubbs, T. F. Cooke, and L. A. Melsheimer, Amer. Ink Maker, 29, 35,425 (1950). ( 5 ) A. A. Ebert, Jr., and H. B. Gottleib, J . Amer. Chem. SOC.,74, 2806 (1952). (6) F.W.Karasek and J. C. Decius, ibid., 74, 4716 (1952). (7) M.Shigemitsu, Bull. Chem. SOC.Jap., 3 2 , 607 (1959). (8) D.N. Kendall, Anal. Chem., 2 5 , 382 (1953). (9) A. N. Sidorov and I. P. Kotlyar, Opt. Spektrosk., 11, 92 (1961). (10) J. M.Assour, J . Phys. Chem., 69, 2295 (1966). (11) J. W. Eastes, U. S. Patent 2,770,620(1956). (12) J. F, Byrne and P. F. Kurz, U. S. Patent 3,357,989(1967); J. H. Sharp and M. Lardon, J . Phys. Chem., 7 2 , 3230 (1968).
Volume 74,Number 0 September 1968
3336
NOTES
and KotlyarQreported that the a form is converted to the P phase when heated above 300". The purpose of this work was to investigate the kinetics of the a+ thermal conversion and the thermal stability of the xpolymorph.
kT h
k~ = - exp(AS*/R) exp(-AH*/RT)
(3)
where k = 1.38 X 10-'6 erg/deg and h = 6.62 X erg sec and is found to be -23.1 f 1.0 cal/mol deg a t 300".
Experimental Section The samples of a-phthalocyanine were prepared by vacuum evaporation of commercially available phthalocyanine onto 1.5- X 0.75- X 0.045-in. rectangular KBr flats obtained from the Perkin-Elmer Corp., Norwalk, Conn. The evaporations were carried out a t torr in a Bendix Balzers Model BA-3 evaporator. The prepared samples were mounted in a simply designed furnace which was placed in the sample beam of a Perkin-Elmer Model 337 grating infrared spectrophotometer. The furnace temperature was varied by adjusting the voltage supplied to a heating tape which was uniformly wrapped about the cylindrical furnace. Three thermocouples were situated radially around the sample, and the temperature of the sample could be maintained to within f2" throughout the temperature range used. Before any conversions were initiated, the samples were held at a temperature of approximately 200" for 8-12 hr.
Results and Discussion The infrared absorption spectrum of the initial a-polymorph in the 700-80O-cm-' region is shown in Figure l a , and the corresponding spectrum of the /3polymorph, obtained by total thermal conversion of the a film, is shown in Figure lb. The kinetics of the a-.P thermal conversion were followed by continuously monitoring the change in the infrared spectra a t a given temperature. The growth of the 724- and 782cm-' frequencies, characteristic of only the /%polymorph, were chosen for the kinetic analyses. If firstorder kinetics are used to describe the thermal conversion, then the rate constant, 161, can be obtained from
12
MICRONS MICRONS 13 14 15 I2 13 14 I5
FREOUENCY (CM'I)
Figure 1. The infrared absorption spectra of metal-free CY- and @-phthalocyaninein the 700-800 em-' absorption region.
26OOC
e
02K/
where A" is the absorbance of the P-polymorph at total conversion and is A t the absorbance at time t. The results, a t five different temperatures between 260 and 345", were plotted according to eq 1, and the first-order plots for the 724-cm-' frequency are shown in Figure 2. The calculated first-order rate constants are listed in Table I and a least-squares analysis of the data fits an Arrhenius plot represented by k1(760 mm) = 108.06*o.08 X exp[(-31,200
370)/RT] sec-'
(2)
' 4,0100 ' &do0 ' 12,000 Time (sec)
'
16LOO
' 20,'oOO' 24&0
Figure 2. The plots of the first-order rate constant calculated from the 724-em-' frequency of the p-polymorph.
Table I : Calculated First-Order Rate Constants Temp,
f
These results correspond to a frequency factor of lo8 sec-' and an activation enthalpy, AH* of 31.2 kcal/ mol. The activation entropy, AS*, of the conversion can be computed from The Journal of Physical Chemistry
O'
OC
260 f.2 298 f 2 303 f 2 320 2 343 f 2
-Rate constant, ki, sec-1724 om-1 782 cm-1
2.12 x 10-6 1.50 X 10-4 2.03 X 4.22 X 1.35 x 10-8
2.04 X 1.38 x 10-4 2.18 X 4.60 X 1.24 X 10-8
(av), ki 880-1
(2.08 f 0.04) (1.44 zk 0.06) (2.10 zk 0.07) (4.41 f. 0.19) (1.30 f.0.06)
X lom6 X X X lo-' X lo-*
NOTES Although the rate data do not yield any information on the mechanism of the conversion process, bond rupture in the phthalocyanine molecule is improbable. The activation enthalpy must be associated with the rearrangement of the phthalocyanine molecules in the CY-.@ polymorphic conversion. The monitored frequencies in the 700-800-~m-~region are associated with the out-,of-plane CH-bending absorption of the four peripherd benzene rings of the phthalocyanine molecule. Since phthalocyanine is a planar molecule, these bending modes will be influenced by the orientation of adjacent molecules which determines the polymorphic form. If the activated complex is similar in structure to the product of the conversion; i.e., @-phthalocyanine,then the negative value of the activation entropy must be associated with an increase in order in the @-polymorph. This is consistent with the fact that single crystals of @-phthalocyanine are easily grown, whereas single crystals of the a-polymorph have not been reported. A close inspection of the infrared spectra resulting from the continuous monitoring of the a+@ conversion gives no evidence that x-phthalocyanine12 is an intermediate. Furthermore, an attempt to thermally convert x-phthalocyanine to another polymorph under conditions similar to those used in the a 4 p conversion was unsuccessful and the x form was found t o be stable up to at least 378”.
Cobalt-60 Radiolysis of Aqueous Eosin1& by A. F. Rodlde, Jr., and L. I. Grossweinerlb Department of Radiation Theram, Michael Reese Hospital and Medical Center, Chicago, Illinois 60616 (Received March 87, 1968)
Organic dyes in aqueous solution are interesting radiolysis systems because of their reactivity with oxidizing and reducing radicals to form strongly colored intermediates which are important also in sensitized photochemical reactions. However, with the exception of methylene blue, limited information on dye radiolysis G values or mechanisms appears in the literature. The only published study on the well-known photosensitizer eosin (tetrabromofluorescein) is the early work of Patti,2 which reports that X-ray irradiation of air-free or air-saturated solutions produces first a colored, nonfluorescent product followed by eventual decoloration. This note reports initial G values for the decoloration of aqueous eosin in the presence of specific scavengers for hydrated electrons, H atoms, and OH radicals. The results are explained by extending the scheme proposed in recent pulse-radiolysis investigations of eosinSaand fluorescein.ab
3337 The irradiation source was a cobalt-60 therapy unit providing a uniform dose rate of approximately 1 rad/ sec in a reaction vessel which could be positioned accurately. The exact dose was calibrated periodically with the modified Fricke dosimeter4 and followed the 5.2-year half-life to within 2%. The irradiation vessel was an 8-ml Pyrex ampoule, provided with a l-cm path side tube for in situ optical absorption measurements, which could be evacuated to a partial air pressure of 2 X torr. A check of the dosimetry procedure with deaerated methylene blue in the presence of 100 mM sodium formate gave G(-MB+) = 2.85 (Table I) in Table I: Initial G ( - S ) Values for Bleaching of Aqueous Eosin by Cobalt-60 y Rays Dye concn,
Q(-W
AiM
PH
10 10 10 10 10 10 10 50 10 (fluorescein) 10 (methylene blue) 10
8.3 8.1 8.1 8.1 8.1 8.2 8.2 8.2 8.1 6.9
0.02 mM HzOz 0 . 2 m M HzOz l.OmMHzOz 18mMNzO 0.25 m M O2 1 m M HCOO100 mM HCOO100mMHCOO100 mM HCOO100mMHCOO-
1.03* (2) 1.43* (2) 1.48* (2) 1.70* (2) 0.85* (2) 1.64 (2) 2.06 (2) 2.08 2.32 2.85
8.2
0.00
10
8.2
1 10 50
8.1 8.1
100mMHCOO+0.25mMO2 100mMHCOO1 mM HzOz Deaerated Deaerated Deaerated
8.7
Additive (5)
+
O.OO(2)
0.33* 0.72* 0.95*
a The asterisk indicates correction was made for growth of colored product; (2) indicates average of two determinations.
agreement with published values of 3.36 and 2.75 (15%)6 and an average of results obtained with different organic reductants of 2.9 f 0.1.’ The samples were prepared with 40 p M NazB407 buffer in triply distilled water. The dye was purified on an alumina-talc column by the method of Koch8 to give emax (518 mp) = (1) (a) Based in part on an M.S. thesis submitted by A. F. Rodde, Jr., to the Physics Department, Illinois Institute of Technology, June 1967. (b) Physics Department, Illinois Institute of Technology, (2) F. Patti, J . Chim. Phys., 52, 77 (1955). (3) (a) J. Chrysochoos, J. Ovadia, and L. I. Grossweiner, J . Phys. Chem., 71, 1629 (1967); (b) P. Cordier and L. I. Grossweiner, ibid., 72, 2018 (1968). (4) L. M. Dorfman and M. S. Matheson, Progr. Reaction Kinetics, 3, 237 (1965). (5) E. Hayon, G. Scholes, and J. J. Weiss, J . Chem. SOC.,301 (1957). ( 6 ) J. P. Keene, E. J. Land, and A. J. Swallow in “Pulse Radiolysis,” M. Ebert, J. P. Keene, A. J. Swallow, and J. H. Baxendale, Ed., Academic Press, Inc., New York, N. Y.,1965,pp 227-245. (7) A. J. Swallow, “Radiation Chemistry of Organic Compounds,” Pergamon Press, Oxford, 1960,pp 175-185.
Volume 78, Number 9 September 1068
