1667
NOTES
L'
! i o ' !so' 'do'
llAl' EXPOSURE TIME, MIN.
'140'
Figure 1. Transfer of C1" from HCl(Cl*B) to AgCl and CdC12as a function of time and temperature. Ratio of C1 atoms in HCl t o C1 atoms in metal chloride is 1:38.
tion of the minimum temperature ranges at which it can be readily observed. The comparisons are qualitative rather than quantitative, since the samples were prepared by different methods and the specific surface areas and defect concentrations were not the same. Except as otherwise noted, the estimated particle sizes were in the 5-p range. A striking feature of the results is the fact that the alkaline earth chlorides all undergo exchange much more rapidly than the alkali metal chlorides, and that their exchange is readily observable at temperatures as low as 0.3 of the absolute melting point. This was true both for samples prepared by dehydration of the hydrates and for crystals of BaC12 and SrClz prepared by cooling from the fused state, to ensure minimum surface are&. The evidence2 indicates that the transfer of chlorine atoms at the surface of the samples of Table I was, in all cases, fast compared to diffusion of the atoms into the latt,ice. The most plausible mechanism for this rapid transfer involves attraction of the chlorine end of an HC1 dipole to a positive lattice site, followed by departure of the proton in combination with a chlorine ion from another site. A mechanism of this type, involving induced dipoles, has been suggested to explain the fact that cc14 exchanges chlorine readily with the ionic surface of aluminum chloride, but not with dissolved or gaseous aluminum chloride.14& The rapid exchange of HC1 with the covalent species SnCL, TiCl,, GaC4, and GeCI4 at -78' suggests that these compounds have ionic surfaces or that a different mechanism, such as complex formation, is available. We have found no significant correlation of the observed exchange rates with available lattice parameters of the chlorides tested.2 Friedel-Craf ts Studies. Studies with AlCld4 and with ZrC146 have shown that they are capable of cata-
lyzing vapor phase Friedel-Crafts reactions as well as exchanging readily with HC1. To determine whether such catalytic effectiveness might be a general property of metal chlorides which exchange chlorine with HCI, the catalytic ability of additional compounds was investigated. These were chosen to include one compound which undergoes extensive exchange with HC1 only at elevated temperature (NaCl), and five compounds representing different types which exchange at room temperature (BaC12, SrCL, ScC4, HfCI4,and PbC12). The NaCl, BaCI2,SrC12,and PbClz were each allowed to stand in contact with ethyl chloride and benzene vapors €or times up to 6 months at 140'. The catalyst was present in large excess relative to the organic reactants which were present in equimolar amounts. The yield of ethylbenzene was a fraction of a per cent at most. ScCb showed no catalytic activity when maintained with the organic materials for 18 hr at 100'. HfCL gave a 30% yield of ethylbenzene in 16 hr at 30'. It is concluded that the ability of an inorganic chloride to exchange chlorine with HCI does not give information about its catalytic efficiency.
Chemical Shifts of Methyl Protons in Methylated Polynuclear Aromatic Hydrocarbons' by I. C. Lewis U n i o n Carbide Corporation, Carbon Products Division, P a r m a Technical Center, P a r m a , Ohio 4.6130 (Received J u l y $0, 1965)
The proton chemical shifts in polynuclear aromatic hydrocarbons have been the subject of considerable interest. The large downfield shifts exhibited by aromatic protons have been attributed to the effects of circula' ing ring currents involving the T electrons.2a Aromztic proton shifts additionally reflect changes in the r-electron density at the attached carbon atom.2b The magnitude of the ring current effect has been accurately estimated for benzene.3 In the treatment of (1) This research was sponsored in part by the Air Force Materials Laboratory, Research and Technology Division, Air Force Systems Command, U. S. Air Force. (2) (a) J. A. Pople, W. G. Schneider, and H. J. Bernstein, "HighResolution Nuclear Magnetic Resonance." McGraw-Hill Book Co., Inc., New York, N. Y., 1959; (b) T. Schaefer and W. G. Schneider, Can. J . Chem., 41, 966 (1963).
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1668
polycyclic hydrocarbons, the ring current intensity for each ring had originally been considered equal to that of benzene. Proton shifts in polynuclear hydrocarbons have been computed from a model based on this principle. 2* Recently, Jonathan, Gordon, and Dailey4 have refined these calculations and computed aromatic proton shifts, u,, for polycyclic hydrocarbons, based on ring current intensities determined individually for each polycyclic system. Only fair quantitative agreement was found between the calculated and measured values of u,. The accurate determination of aromatic proton chemical shifts is complicated by spin-spin splittings. Furthermore, aromatic proton shifts are quite sensitive to nonuniform solvent effects6 and to steric effects.*s6 These factors present limitations to a quantitative test of the ring current model in polynuclear hydrocarbons. The ring current model can alternately be tested by the nmr absorptions of aromatic methyl protons which can be measured precisely and unambiguously. I n addition, the protons in methylated aromatics have been found to be less susceptible to solvent effects than those of unsubstituted aromatics.'~~ Maclean and Mackorg measured the nmr spectra for a limited number of methylated aromatics at 40 Mc/sec and showed a relationship between the methyl proton shifts and those of the aromatic protons in the unsubstituted hydrocarbon. Durand, Parello, and Buu-HoilO have measured the methyl proton shifts for additional methylated polynuclear aromatics a t 60 hlc/sec. These investigators found a relatively poor agreement between the experimental parameters and those computed from the ring current model in which the current intensity of benzene was employed for each ring unit. The present investigation reports methyl proton chemical shifts which have been determined at 60 R/lc/sec for a variety of polycyclic methyl compounds. These quantities are shown to agree quite well with the refined u, values determined by Jonathan, et aL4
Experimental Section The aromatic hydrocarbons employed in this study, with the excbeption of methylcoronene, were obtained from commercial sources and generally used as received. I n :% number of instances where the melting behavior indicated some impurity, the compounds were crystallized from organic solvents. Methylcoronene was synthesized by the Friedel-Craft methylation of coronene." The solutions for measurement were made up to 5% of hydrocarbon in CC1, and contained (CH&Si as an internal standard. The spectrum for methylcoronene The Journal of l'hysical Chemistry
NOTES
Table I: Chemical on Aromatic Rings Ring system
Benzene
Shifts of Methyl Groups Substituted
Methyl substituent
1
0.00
2.68 2.50 2.47, 2.57 2.70 2.50, 2.63 2.40 2.38, 2.47
0.85 0.42
1 2 3 9
2.73 2.53 2.62 2.72"
0.99 0.56 0.61 0.99
2 9 9, 10
2.55 3.05 3.05
0.55 1.83
1 2 4
2.93 2.80 2.87
1.55 1.12 1.40
1
3.3@
2.74
11 3 1, 4 1, 3, 1, 2, 1, 2, 1, 2,
:&
1, 2
7&2
Anthracene
Jfc& 6
4
10
Pyrene
6
5 4, 5 3, 4, 5 3, 4, 5, 6
1 2 1, 1, 21 2,
Phenanthrene
7r
2.37 2.25 2.30 2.28 2.25 2.15 2.13, 2.18 2.17
1, 2
Naphthalene
8, ppm
5 6 3 3, 6
5
Coronene
@ Data of ref 8.
' Measured in CSa with reference to toluene.
(3) J. S. Waugh and R. W. Fessenden, J . Am. Chem. SOC.,79, 849 (1957). (4) N. Jonathan, S. Gordon, and B. P. Dailey, J . Chem. Phys., 36, 2443 (1962). (5) A. A. Bothner-By and R. E. Click, ibid., 26, 1651 (1957). (6) C . Reid, J. Mol. Spedry., 1,18 (1957). (7) See ref 1, p 426. (8) F. F. Yew, R. J. Kurland, and B. J. Mair, Anal. Chem., 36, 843 (1964). (9) C.Maclean and E. L. Mackor, Mol. Phys., 4,241 (1961). (10) P. Durand, J. Parello, and N. P. Buu-Hoi, Bull. SOC.Chim. France, 2438 (1963). (11) The author wishes to thank Mrs. 6. Wallon for the synthesis of methylcoronene.
NOTES
1669
-
3.0
2.5
2.0 cr I .5
the relative contribution of the circulating ring currents to the chemical shifts of the aromatic ring protons and the aromatic methyl protons. The magnitude of this slope can be estimated by employing the model of Johnson and Bovey12to calculate the ring current contributions to the aromatic methyl shifts.1° A comparison of these computed methyl shifts with those calculated for the ring protons4 leads to a slope of 2.1 which is in good agreement with the slope of 2.5 for the experimental methyl shift data in Figure 1. Several polymethylated aromatics are also included in Table I. I n general, additional methyl substitution in the ring shifts the methyl proton resonance upfield. This result is expected in view of the electron-releasing nature of the methyl group and has been described elsewhere.4 (12) C. E. Johnson and F. A. Bovey, J . Chem. P h y s . , 29, 1012 (1958).
Changes in Dielectric Relaxation during Dehydration and Rehydration of Rochelle Salt Figrire 1. Relationship between computed aromatic proton shifts, uT. and methyl proton shifts, 6, in polycyclic hydrocarbons.
by P. G. Hall and F. C. Tompkins Department of Chemistry, Imperial College of Science and Technology, London, S . W.7, England (Received September 27, 1965)
was obtained in CSZ and measured relative to toluene in the same solvent. Spectra were recorded at 60 Rlc/sec with a Varian 4-60 spectrometer.
Results Summarized in Table I are the chemical shifts, 6, in ppm for the protons of methyl groups substituted in various positions of six aromatic ring systems : benzene. naphthalene, phenanthrene, anthracene, pyrene, and coronene. Table I also lists the screening constants, uT, for the aromatic protons at the corresponding positions in the ring as calculated by Jonathan, et aL4 Figure 1 illustrates a comparison of the 6 CH? values with the uT aromatic proton shifts. The linear agreement is quite good and provides further validation of the ring current model as well as of the conclusions of Maclean and 3Iackor pertaining to the relationship between aromatic methyl and aromatic proton chemical shifts. The ninr methyl data in Table I show a much poorer correlai,ion with the aromatic proton chemical shifts computed from the single benzene ring current model.2a The slope of the line in Figure 1 represents
Garner1 has reviewed early gravimetric studies of the rehydration of dehydrated salt hydrates. Discontinuities in the plots of water uptake against time did not correspond to any known hydrates. With common alurnl2multilayer adsorbed water controlled the rate of diffusion into the porous anhydride, but a phase transformation corresponding to A1,(SOS3.9.4H20 produced a “glassy” modification which was very stable to further rehydration after dehydration. JIore recent investigations, also using gravimetric methods, with lead styphnate3 and manganous oxalate4 showed that the conditions of dehydration can have a pronounced effect on the subsequent rehydration. This note concerns the use of a dielectric relaxation (1) W. E. Garner, E d , “Chemlstry of the Solid State,” Butterworth and Co. (Publishers) Ltd., London, 1955, Chapter 8 (2) A. Bielanski and F. C. Tompkins, T r a n s . Faraday SOC.,46, 1072 (1950). (3) T. B. Flanagan, $bid., 5 5 , 114 (1959). (4) T. B. Flanagan and M. K. Goldstein, J . P h y s . Chem., 68, 663 (1964).
Volume 70, Number 5
M a y 1966
