3334 to be 24-25 dyn/cm which is much lower than those of polyethylene and poly(viny1 alcohol) (37 dyn/cm) . H e thus proposed that the CH3 groups appeared to predominat’e over CH2, ether oxygen, and OH groups on the poly(viny1 butyral) surface. However, the yo value calculated by our eq 8 for a poly(viny1 butyral) containing 35 mol % ’ of poly(viny1 alcohol) and having a density of l . O V is 25 dyn/cm, in good agreement with the experimental value, indicating that the various groups areprobably actuallyrandomly distributed on the surface. The low yo value of the polymer as compared with those of polyethylene and poly(viny1 alcohol) arises from the presence of three CH groups in the vinyl butyral segment which lower the molar attraction constant. (13) H.R. Simonds and J. M. Church, “A Concise Guide to Plastics,” Reinhold Publishing Corp., New York, N. Y., 1963,p 70.
Absorption Spectra of Sodium-Ammonia Mixtures in the Gas Phase
by Irving Warshawsky Lewis Research Center, National Aeronautics and Space Administration, Cleveland, Ohio 44156 (Received March 6 , 1968)
Gibson and Argo’s spectral studies of dilute solutions of alkali metals in liquid ammonia suggest that a common species is present in these solutions.1 This species has been identified as the “solvated electron.” In a more recent investigation, where both the metal concentration and spectral ranges were extended, it was concluded that for the solvated electron Beer’s law holds, even at concentrations where ion-pair formation was occurring.2 Thus the position and shape of the absorption band of the solvated electron are surprisingly insensitive to the changes occurring in the electron’s environment. It is of interest to investigate whether the electron from sodium is delocalized in gaseous ammonia as it is when sodium dissolves in liquid ammonia, in view of Naiditch’s demonstration that sodium is soluble in gaseous ammonia. Further, the gaseous medium allows a broad range of relative concentrations of the solvent through the control of pressure. The purpose of this note is to report some observations of the absorption spectra of gaseous sodiumammonia mixtures over a range of ammonia pressures from 0 to 6 atm. A diagram of the optical cell used for the spectral measurements is shown in Figure 1, with the path length between inner windows being 5.00 cm. A trace heater element (consisting of a metal-sheathed oxide-insulated The Journal of Physical Chemistry
NOTES Vac line
bd
Heating element
Heating Constriction element
1
II
I -Ltght path
Ampoule (to condense NH31 Constridton
w Na
Figure 1. Optical cell used for spectral measurements.
nichrome wire) was wrapped around the inner windows of the vacuum-jacketed wall. The main heat input to the cell was by means of silver-coated copper shavings in thermal contact with the inner walls of a small oven. Temperature regulation was obtained with an Electromax pyrometric controller and an iron-constantan thermocouple which was welded to a copper plate placed in the oven. At 250”, temperature fluctuations near the cell were less than 3”. The outer windows of the optical cell were cooled by flushing with nitrogen gas. A Perkin-Elmer 350 recording spectrophotometer was used to record the spectra. In a typical run, the optical cell was glass-blown to the vacuum line and pumped on for approximately 1 week a t pressures less than lob6 mm. During the final 24 hr of pumping, a liquid nitrogen bath was placed around a trap between the McLeod gauge and the cell in order to minimize mercury contamination of the cell due to its possible catalytic a ~ t i v i t y . ~During the distillation of sodium into the cell through the constriction (Figure l),the inner windows were coated with sodium. This sodium was distilled from the inner windows to the walls of the cell by passing a current through the heating element (Figure 1). The excess sodium was sealed off at the constriction while pumping. A known quantity of dry ammonia was then distilled into the cell. This was done by use of a gas buret (ammonia vapor a t a pressure of 4.20 cm was distilled into an ampoule on the vacuum line whose volume had previously been calibrated) along with a liquid nitrogen bath which was placed around the ampoule on the cell shown in Figure 1. With a liquid nitrogen bath around the ampoule, the system was (1) G. E. Gibson and \V. L. Argo, J . Amer. Chem. Soc., 40, 1327 (1918). (2) M.Gold and W. L. Jolly, Inorg. Chem., 1, 818 (1962);M.Gold, W. L. Jolly, and K. S. Pitzer, J . Amer. Chem. SOC.,84, 2264 (1962). (3) 9. Naiditch in “Metal Ammonia Solutions: Physical-Chemical Properties,” G. LePoutre and M. J. Sienko, Ed., W. A. Benjamin, Inc., New York, N. Y., 1964,p 113. (4) J. F. Dewald and G. Lepoutre, J . Amer. Chem. SOC., 76, 3369 (1954).
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
