NOTES
3066
tion with recent discussions of proton chemical shifts of metal ion-DMF c o m p l e ~ e s . ~ ~It- 'has ~ been suggested18 that since formyl proton shifts of DMF complexes of in D M F solutions decrease in the order A1C13 > AlBra > A&, the stability of the complexes might decrease in the same order. As we have shown, the shifts are due not to a first but to a second coordination sphere interaction between Al(DMF)63+and X-. It is difficult to discern the effect that such an interaction would have on the Al(III)-DAIF bond with a variation of X-. It has also been reported18 that the rates of DMF exchange for the DMF complexes of AlX3 in DMF decrease in the order C1 > Br, I. Such an effect, if true, must reside in second coordination sphere interactions. Considering the importance of such an observation in the light of the SN1 mechanism proposed for solvent exchange,2oit would be of interest to reevaluate those rates of DMF exchange using standard line width or complete line shape analysis techniques rather than the peak-height technique which was used.21 A correlation between the formyl proton chemical shifts of a series of metal ion-DMF complexes and the metal ion [charge]: [radiusI2 ratio has been devised.Ig That correlation may be an "accidental" one, because the A1C13complex conforms to it and a large fraction of the formyl proton shift for the A ~ ( D A I F ) Eion ~ + in the presence of C1- is due to the effect of C1- not Al(II1). The BeCl2-DlWF complex also conforms to the correlation, but we have found that, in anhydrous D M F solutions of BeC12,appreciable concentrations of B e c k (DMF)2 and Be(DRilF)42+ exist in equilibrium with [Be(DP\iIF)42+]C1-ion pairs.22 Experimental Section Reagents. Eastman White Label DMF was purified as described previously. 2o Preparation of Complexes. The complex Al(D;\IF) 6was prepared as described previously." The complexes A1(DMF)6XB(X is C1, Br, or I) were prepared by dissolving the appropriate anhydrous A1X3 in DMF, flooding the solid out of solution with diethyl ether, and drying the solid in vacuo at 25". Anal. Calcd for Al(DMF)&l3: Al, 4.73%; C1, 18.6%. Found: Al, 4.73; C1, 18.5. Calcd for Al(Di\lF)GBr3: Al, 3.837,; Br, 35.4%; N, 11.9%. (17) A. Fratiello, D. P. Miller, and R. Schuster, iMol. Phys., 12, 111 (1967). (18) A. Fratiello and R. Schuster, J . Phys. Chem., 71, 1948 (1967). (19) A. Fratiello, R. Schuster, and D. P. AMiller,M o l . Phys., 11, 597 (1966). (20) See, for example, M. Eigen and R. G. Wilkens, "Mechanisms of Inorganic Reactions," American Chemical Society, Washington, D. C., 1965, pp 55-56, and references therein. (21) We have found the line-width technique and especially the complete line-shape analysis more reliable than the method based on peak heights. (22) W. G. Movius and N. A. Matwiyoff, to be submitted for publication. The Journal of Physical Chemistry
Found: Al, 3.70; Br, 33.3; N, 11.9. Calcd for Al(DMF)G13:Al, 3.19%; I, 45.0%. Found: AI, 3.12; I, 43.2. Measurements. Pmr spectra were obtained at 60 MHz using the Varian A-60-A spectrometer. Aluminum-27 nmr spectra were obtained at 12 MHz using the Varian HR-40 spectrometer. The systems were calibrated and the measurements were made in the manner described p r e v i o ~ s l y . ~ Computer ~J~ programs were run on the IBM 08360/67. (23) N. A. Matwiyoff, Inorg. Chem., 5, 788 (1966).
Spectroscopic Evidence for Bransted Acidity in Partially Dehydrated Group Ia Forms of Zeolites X and Y by Yoshihiro Watanabe2 and Henry W. Habgood Research Council of Alberta, Edmonton, Alberta, Canada (Received March 11, 1968)
Recent reports by Liengme and Hall,3 Hughes and White,* Ward,5 and Eberly6 have described investigations of zeolite acidity using infrared spectroscopy of adsorbed pyridine according to the methods developed by Parry7 and Ba~ila.8,~These authors have shown that the acidity of HY zeolite is principally Brpinsted acidity which is converted to Lewis acidity on dehydroxylation, while the acidity of the cationic forms of Y zeolite is principally Lewis or pseudo-LewisGacidity associated with the exchangeable cations. The group I I a forms of Y zeolites5J seem also to have some inherent Brpinsted acidity, which is enhanced by the presence of small amounts of water. The group I a zeolites show no Brgnsted acidity in the dry form, and Wardlo has also found no detectable Brpinsted acidity in slightly hydrated group I a Y zeolites. Experiments that we have been carrying out do, on the other hand, show significant Brpinsted acidity in some samples of (1) Contribution No. 414 from the Research Council of Alberta, Edmonton, Alberta, Canada. (2) RCA Postdoctoral Fellow 1965-1966. (3) B. V. Liengme and W. K. Hall, Trans. Faraday Soc., 62, 3229 (1966). (4) T. R . Hughes and H. M. White, 5.Phys. Chem., 71, 2192 (1967). (5) J. W. Ward, Abstracts, 154th National Meeting of the American Chemical Society, Chicago, Ill., 1967, No. 143; J . Catal., 9, 225 (1967). (6) P. E. Eberly, Jr., Abstracts, 154th National Meeting of the American Chemical Society, Chicago, Ill., 1967, No. 142. (7) E. P. Parry, J . Catal., 2 , 371 (1963). (8) M. R. Basila, T. R. Kantner, and K. H. Rhee, J . Phys. Chem., 68, 3197 (1964). (9) M. R. Basila and T. R. Kantner, ibid., 70, 1681 (1966). (IO) J. W.Ward, J . Catal., in press.
NOTES group Ia zeolites X and Y under conditions of partial dehydration. The details of these observations and some comments on their possible significance are given below. Experimental Section Measurements were made on the Li, Ne, and K forms of zeolites X and Y prepared from Linde Lots 1340080 and SK40 by treatment with 1 N solutions of the appropriate chlorides followed by washing with water made basic to pH 10 with the corresponding hydroxide. The degrees of exchange attained were: LiX, 85%; LiY, 67%; K X and KY, over 95%. The parent NaX and NaY zeolites as received were both cation deficient (Na:AI ratios of 0.97 and 0.89, respectively), and from the analyses of samples from these batches that received similar chloride-hydroxide treatments," we would expect the cation deficiency to be reduced in the Y zeolites but to be increased in the X zeolites. Measurements were also made on a number of other samples of NaX and N a y : the parent NaX (Lot 1340080) and NaY (Lot SK40) as received; NaX (Lot 1340080) after leaching with distilled water or with p H 10 NaOH solution; NaX (Linde Lot 13420, Na:A1 ratio of 0.97), and (Linde Lot 13916, Na:A1 ratio of 0.94) as received; and NaY (Linde Lot 11007-73, a high-purity material with an Na: AI ratio of 0.99 and 0.97 in separate determinations) both as received and after the NaC1NaOH treatment described above. I n the case of NaX leached with pH 10 NaOH, it was confirmed that approximately 5% additional hydrolysis occurred. A sample of NHdY was also prepared in order to obtain HY for a reference spectrum. The spectroscopic measurements were carried out on unsupported pellets of density 10-20 mg/cm2, 0.75 in. in diameter, and pressed at 40,000 psi. The cell has been previously described12 and all spectra were run a t room temperature on a Beckman IR-12 spectrophotometer. Fully dehydrated zeolites were produced by heating under vacuum a t 500" for from 4 to 16 hr. Partial dehydration was achieved by heating the original pellet under vacuum a t temperatures between 150 and 200" for 4-16 hr, and partial hydration was achieved by the addition of known doses of water to the pellet a t 100" following full dehydration a t 500". In most cases the sample was equilibrated with pyridine from a reservoir at 0" and the excess pyridine was pumped off a t 150-200". I n a few experiments, mostly with dry zeolite, known doses of pyridine were added to the sample at 150"; these were completely adsorbed and the integrated molar absorptivities were determined from the spectra. The pyridine showed a considerable tendency to dissolve in the silicone stopcock grease so that the exact quantitative transfer to the pellet was somewhat uncertain. The spectroscopic slit width waa 0.25 or less of the
3067 apparent half-band widths. The maximum band absorbance and the width a t half-height (in terms of absorbance) were measured, and the integrated band intensity was taken as being proportional to their product. Absolute integrated band intensities (for comparison with Hughes and White) were taken as a/2 times this product.la The frequency scale of the spectrometer was calibrated from time to time against the spectrum of water vapor, and the frequencies reported are considered accurate to It 1 cm-'. Results The three characteristic bands in the absorption spectrum of adsorbed pyridine that are of interest in this study are found near 1445, 1490, and 1545 cm-I. Under the fairly severe outgassing conditions used here (pumping a t 150"), the band near 1445 cm-' is considered to represent pyridine strongly adsorbed on a cation site. For convenience we will refer to this as Lewis pyridine, although we recognize that relative to the interaction with the usually accepted trigonal aluminum Lewis site, the cation may better be thought of as a pseudo-Lewis &tee6 The band near 1545 cm-l is characteristic of pyridinium ion and hence represents Brdnsted pyridine, while the band near 1490 cm-l is due to both species. The observed frequencies for these various bands are given in Table I.
Table I: Observed Frequencies (in em-1) for the Principal Absorption Bands of Pyridine Adsorbed on Lewis and Br$nsted Sites of Group Ia Zeolites Frequenoy, cm-l-
I--
Lewis sites Lewis-BrZnsted sites Brpsted sites
7
LiX
NaX
KX
LiY
NaY
KY
1443 1492
1445 1491
1441 1489
1446 1491
1446 1491
1442 1489
1544
1545
1549
1545
1545
1546
The integrated molar absorptivities for these bands are shown in Table I1 and are seen to be in reasonable agreement with the more extensive measurements of Hughes and White. Although one might reasonably expect some variation in the molar intensities of, particularly, Lewis bands with different adsorbents, the measurements are not suffciently precise to establish this, and a single average value is taken as a first approximation. These results do confirm the quite wide difference in intensity of the 1490-cm-1 band in Lewis(11) H. W. Habgood and Z. M. George, "Proceedings of the S.C.I. Conference on Molecular Sieves, London, 1957." (12) L. Bertmh and H. W. Habgood, J . Phys. Chem., 67, 1621 (1 963). (13) R. N. Jones and C. Sandorfy in "Techniques of Organic Chemistry,'' Vol. IX, A. Weissberger, Ed., Interscience Publishers, New York, N. Y., 1966,p 284.
Volume 78, Number 8 August 1968
3068
NOTES ~-
Table I1 : Integrated Molar Absorptivities of Characteristic Brplnsted and Lewis Absorption Bands of Adsorbed Pyridine Apparent integrated molar absorption intensities (adsorbent a8 shown), Molecular species
Absorption freguency,
This
cm-1
work
PyB" PYB PyLb
1490 1545 1490
PYL
1445
cm/fimol-
2 . 4 (HY) 2 . 5 (HY) 0.72 (LiX) 0 . 5 8 (NaX) 0.91 (NaX)" 0.65 (Nay) 2 . 4 (LiX) 3 . 1 (NaX)
Hughes and White
3 . 4 (HY) 3 . 0 (HY) 0.56 (alumina)
3 . 3 (alumina) 3 . 9 (silicaalumina)
4 . 8 (NaX)" 4 . 1 (Nay)
'
a Brplnsted-bound pyridine. Lewis-bound pyridine. ' A second sample from the same lot but measured in a different infrared cell.
bound pyridine and the same band in the pyridinium ion. I n agreement with Ward5Jo and Eberly,'j no detectable band a t 1545 cm-l was seen in the spectrum of adsorbed pyridine on any of the dehydrated group I a zeolites. Neither was any Brginsted band seen when pyridine was adsorbed on partially dehydrated NaX taken directly as received or on any of the zeolites under conditions of partial hydration. However, on the samples of X zeolite that had been subjected to any of the aqueous treatments described above and on all samples of Y zeolite, the addition of pyridine after partial dehydration resulted in the appearance of the pyridinium absorption band at 1545 cm-'. The development of maximum intensity depended somewhat upon the exact conditions of dehydration. For example, dehydration with pumping at 150" for a few hours followed by adsorption of excess pyridine and pumping at 150" for 3-4 hr gave a well-developed band a t 1545 cm-l; further pumping at 150-200" frequently increased the intensity of this band. Taking the pellet through a cycle of dehydration at 500°, complete rehydration, partial dehydration, and adsorption of pyridine a t 150-200" also gave the 1545-cm-' band for pyridinium ion. With some X zeolites, the intensity of the 1545-cm-1 band following the second partial dehydration was lower than that corresponding t o the initial partial dehydration, suggesting some degree of irreversible behavior. One set of measurements obtained with a series of X and Y zeolites partially dehydrated by pumping overnight at 170" is shown in Table 111. The results suggest a decrease in the concentration of Brgnsted sites and an increase in the concentration of Lewis sites in the series Li, Na, and K, with the total number of sites The Journal of Physical Chemistry
remaining roughly constant at about 0.5 site/large cavity in the X zeolites and 0.2 site/large cavity in the Y zeolites. A more limited series of measurements on the very pure NaY zeolite gave similar results with, if anything, even higher concentrations of Brginsted sites (125 and 190 mequiv/equiv of large cavities for Lewis and Brginsted sites in Nay). The NaX zeolites in their as-received form showed infrared absorption bands corresponding to chemisorbed carbon dioxide12 and, as mentioned above, did not develop Brginsted acidity on partial dehydration. Complete dehydration to 500", complete rehydration with water vapor, and a second partial dehydration substantially removed the extraneous infrared bands but did not lead to Brginsted acidity. Apparently a more intensive aqueous treatment is required for this. Table 111: Concentrations of Lewis and Brplnsted Acid Sites on Partially Dehydrated Zeolites ---Conon, mequiv/equiv of large oavities--7 LiX NaX KX LiY NaY KY
Lewis sites" Brplnsted sites* Ratio Lewis/ Br@nsted
280 140 2.0
350 130 2.7
410 52 7.9
86 90
1.0
130 60 2.2
140 40 3.5
a From the 1445-cm-l band of adsorbed pyridine; the inteFrom the grated molar absorptivity taken as 3.6 cm/pmol. 1545-cm-1 band of adsorbed pyridine; the integrated molar absorptivity taken as 2.5 cm/pmol.
'
Examination of the OH-stretching region in the infrared spectrum showed the presence of sharp nonhydrogen-bonded OH-stretching absorption bands characteristic of small amounts of adsorbed water12in both the partially dehydrated and the partially hydrated zeolites, and there was no change in the intensities of these bands when small amounts of pyridine were added. The various cationic forms of the X zeolite from Lot 1340080 showed significant absorption a t 3650 cm-l, suggesting some partial hydrolysis and generation of H X sites, but, in contrast with the behavior reported for HY,4 this absorption band likewise was unchanged when the Brgnsted interaction with pyridine occurred. The only positive observation of a change in the OH or water absorption spectrum accompanying the formation of pyridinium ion was a decrease in the HOH-bending band a t 1650 cm-l and a broad increase on the low-frequency flank of the hydrogenbonded OH-stretching band, roughly over the range 3000-2400 cm-l. We had expected that the Brginsted acidity might be associated with the trace amounts of water adsorbed by the cations in such manner as t o leave one hydroxyl group free of hydrogen bonding.12 The proton of this
'NOTES free OH should be decidedly more acidic than normal protons in water, but apparently it is not acidic enough to react with pyridine. Water and pyridine added in low concentrations to the dehydrated zeolites adsorb on separate cations as indicated by the simultaneous presence of the sharp OH band (3695 cm-l for NaX) of adsorbed water and the pseudo-Lewis band of pyridine adsorbed on the cation (1445 cm-l on NaX). Alternatively, the Brgnsted acidity might be due to hydrogen sites formed by partial hydrolysis of the cations and hence be a measure of the cation deficiency. A 1% cation deficiency corresponds to around 110 mequiv/equiv of large cavities in X zeolite and to around 70 mequiv/equiv of large cavities in Y zeolite. With the possible exception of the high-purity N a y zeolite, the observed cation deficiencies are all more than sufficient to account for the Brgnsted acidity, and even for the high-purity NaY the 3y0 cation deficiency of the second analysis would correspond to 210 mequiv/ equiv of large cavities, as compared with 190 mequiv/ equiv of large cavities observed. On the other hand, the X zeolites which were already cation deficient as received did not show Brgnsted acidity until they had undergone some further hydrolysis. Also, although the Brgnsted sites disappeared upon heating to 500°, they were not replaced by Lewis sites (presumably exposed aluminum ions) that with adsorbed pyridine give a band at 1452 cm-' (easily distinguishable from the 1441-1446-cm-' band) as is the case for HY.4 The possible role of cation deficiences could be clarified by experiments with a carefully prepared zeolite sample having the stoichiometric concentration of cations. The development of Brgnsted acidity by group I a zeolites in the presence of small amounts of water had been indicated by earlier kinetic studies of cyclopropane i s ~ m e r i z a t i o n , ~although ~J~ the results of Habgood and George1' suggested that concentrations of active centers were probably much lower than the concentrations of Br@nstedsites given in Table 111. The results, taken as a whole, indicate that the Br@nsted acidity present on these group I a zeolites is associated with some sort of metastable structural element in the zeolite crystal-one which disappears on intense dehydration but reforms after complete rehydration. While these sites may in some way be associated with cation deficiency, it is tempting to look for an association of these effects with the reported shifts in cation positions16or with the apparent change from octahedral to tetrahedral coordination around some aluminum ions16accompanying dehydration. It is also interesting that the observed concentrations of acid sites given in Table I11 (0.2-0.5 site/cavity) are similar to the concentrations of various other cation-associated strong adsorptions found in the group I a zeo1ites.l' Acknowledgment. We thank John W. Ward of Union Oil Co., Brea, Calif., for several helpful discussions and
3069 for making available some of his results prior to publication. (14) D. W. Bassett and H. W. Habgood, J. Phys. Chem., 64, 769 (1960). (16) K. Seff and D . P. Shoemaker, Acta Crystallogr., 22, 162 (1967). (16) J. Turkevich, Catal. Rev., 1, 1 (1967). (17) H. W. Habgood, Chem. Eng. Progr. Symp. Ser., 63, No. 73, 46 (1967).
Observation of a Minimum in the Kerr Constants of Light and Heavy Water Near 30"' by Yeong-jgi Chen and William H. Orttung Department of Chemistry, University of California, Riverside, California 8.9601 (Received April 3,1968)
Interpretation of the Kerr effect of water has become a problem of considerable interest in recent years. A general theory of the Kerr constant of liquids was developed and applied to water by Buckingham and RaabS2 Experiments to test this theory yielded data a t visible wavelengths and 25°.3 Waibe14has recently reported dispersion data from the visible wavelength region to 2000 A, Data on the temperature dependence of the Kerr constants of H2O and D2O from 5 to 55" a t 365 mp (and also a t 436 mp for D2O) are reported in this note. Studies of the wavelength dependence should yield information about the electronic and vibrational transitions that,contribute to the Kerr effect. . The temperature dependence should help to decide the relative importance of effects such as hyperpolarizability (change of polarizability with field strength), induced-dipole orientation, and permanent-dipole orientation. The temperature dependence may also reflect the nature of the orientational correlations of neighboring molecules.
Experimental Section The optical part of the apparatus consisted of a light source, monochromator, beam splitter, polarizer, cell, analyzer, and photomultipliers. A 200-W mercury arc Iamp was used, followed by an Engis f/lO monochromator. The polarizer and analyzer were glan prisms from the Crystal Optics Co. The cell had a 10mm path length and used microscope slide windows. The dipping electrodes had a separation of 1.47 mm. (1) This investigation was supported in part by Public Health Service Grant GM11683 from the Division of General Medical Sciences. (2) A. D. Buckingham and R . E. Raab, J . Chem. SOC.,2341 (1957). (3) W. H. Orttung and J. A. Meyers, J. Phys. Chem., 67, 1906 (1963). (4) J. Waibel, 2. Naturforsch., 21a, 186 (1966). Volume 7.9,Number 8
August 1968
