3063
NOTES
result in a more proficient rupture of the oyclobutane moiety of the parent ion in the order I > I11 > 11. Indeed, the relative abundance of the C7H7+ ion which requires the cyclobutane to remain intact (reaction 2) is in the expected inverse order I1 > I11 > I.
ENW-TRANS-ENDO
Magnetic Resonance Studies of Outer-Sphere Ion-Pair Formation in N,N-Dimethylformamide Solutions of Aluminum(II1) Halides 4 2 0
20 40
I I h . J , do. I l d o
do do
Id0
I I]
by W. G. Movius and N. A. Matwiyoff
Id0 Id0 200
Chemistry Department, Pennsylvania State University, University Park, Pennsylvania 16802 (Received March 18, 1968)
lll/b
Figure 3. Fragmentation of pentacyclotetradecane stereoisomers by 15-eV electrons.
to be the average of compounds I and 11. While the spectrum of compound I11 is indeed intermediate, it is not an algebraic average, which suggests that nonbonded effectsa are involved. Table I : Fragmentation of Penetacyclodecanes Using 12-eV Electrons d e
ezo,trans,ezo I
endo,trans,endo I1
endo,trans,ezo I11
159.0 160.0 188.0 189.0
14.21 5.68 68.75 11.36
49.00 16.33 29.65 5.03
28.57 17.23 46.64 7.56
Several factors could be responsible for this observation. The ease of hydrogen atom transfer accompanying formation of many dominant fragment ions may be steric dependent, but the large difference in parent ion abundance does not reflect itself in a corresponding change in the intensities of the It-28 and M-29 peaks even a t 12-eV nominal ionizing voltage where successive dissociation is minimized (Table I). Small variations in intramolecular strain exist in these isomers and could be of influence, but this is difficult to assess. Another attractive explanation is based on the enhanced stability of norbornal ions with the leaving groups in the exo oonfigurationl7 as suggested for the somewhat analogous case of the bromonorbornanes.s This would (6) R. I. Reed, “Ion Production by Electron Impact,” Academic Press Inc., New York, N. Y., 1962; “Application of Mass Spectrometry to Organic Chemistry,” Academic Press Inc., New York, N. Y., 1966. (7) For a recent review of this area, see P. D. Bartlett, “Nonclassical Ions: Reprints and Commentary,” W. A. Benjamin, Inc., New York, N. Y . , 1965. (8) D. C. DeJongh and S. R. Shrader, J. Amer. Chem. Soc., 88, 3881 (1966).
It is, in general, not possible to distinguish between intimate and solvent-separated ion pairs with the use of common activity methods for probing ion-pair formation in electrolyte solutions, e.g., colligative conductivity, or electromotive force measurements.1~2For transition metal ions having well-defined first coordination spheres, such a distinction, in certain cases, can be made with the aid of proton magnetic resonance spectroscopy,6-7 e.g., to distinguish ( ~ m r ) ~or> optical * { [Co(NH&OHJS04\ (as) from [c0(NH&S0d+(aq).~ For a few diamagnetic cations having doand dl0electronic configurations, intimate ion pairs have been detected by using pmrs~eand Ramanlo spectroscopy. We wish to report here, for solutions of aluminum halides (AlXB) in N,N-dimethylformamide (DMF), pmr data which can be interpreted in terms of a second coordination sphere interaction between a well-defined A1(DMF)e3+ ion and X-. The pmr spectra of anhydrous Al(C104)a in DMF consist of two sets of signals, one set due to the formyl and nonequivalent N-methyl signals of DMF coordinated in the ion A1(DR/IF)63+and the other set due to anal+
(1) W. J. Hamer, Ed., “The Structure of Electrolyte Solutions,” John Wiley and Sons, Inc., New York, N. Y., 1959. (2) C. B. Monk, “Electrolytic Dissociation,” Academic Press Inc., New York, N. Y., 1961. (3) Z. Luz, J . Chem. Phys., 41, 1748 (1964). (4) B. M. Fung, J. Amer. Chem. Soc., 89, 5788 (1967). (5) F. A. Posey and H. Taube, ibid., 78, 15 (1956). (6) C. H. Langford and W. R. Muir, J. Phys. Chem., 71, 2602 (1967). (7) C. H.Langford and W. R. Muir, J . Amer. Chem. Soc., 89, 3141 (1967),and references therein. (8) N. A. Matwiyoff and W. G. Movius, ibid., 89, 6077 (1967). (9) W. G. Movius and N. A. Matwiyoff, Abstracts, 3rd Middle Atlantic Meeting of the American Chemical Society, Philadelphia, Pa., Feb 1968,p 34. (10) D. E. Irish in “Raman Spectroscopy,” H. A. Szymanski, Ed., Plenum Press, Ino., New York, N. Y.,1967. Volume 78, Number 8 August 1968
3064
NOTES
ogous signals of DMF in the bulk of the solution." The chemical shifts of these signals are independent of the concentration of Al(C10& (0.05-0.2 m). Upon the addition of LiC104 (up to 1.0 m) to these solutions, the chemical shifts of the A1(DMF)03+protons do not change and the proton signals of bulk DMF shift only slightly (a maximum of 3 Hz downfield for the formyl proton). In contrast to the behavior of A1(C104)3,DRIF solutions of AlXa exhibit large chemical shifts of the formyl proton resonance of coordinated DMF as a function of X- ion concentration. The chemical shifts of the protons in 0.1 m DMF solutions of A1X3 are summarized in Table I. The protons corresponding to the assignments in Table I are defined in Figure 1. The difference in the shifts between the A1(C104)3 and AlXa solutions must arise from a second coordination sphere interaction, because the relative areas of the coordinated and bulk DMF signals are consistent, in each case, with a coordination number of 6.0 rt 0.1 for Al(II1); that is, the predominant Al(II1) complex in each solution is A1(DMF)e3+. A downfield shift in the protons of coordinated DRIF can also be induced by adding anhydrous LiX to DR4F solutions of Al(ClO4)s. As noted above, addition of LiC104to solutions of Al(C104)a caused no shift of the coordinated and only slight shifts in the bulk DMF signals.
lb')
Free DMF
Coordinated D M F
Figure 1. DMF proton sites corresponding to the assignments in Table I.
could be extended into the range 15-80, and a limiting value of the shift, independent of X- ion concentration, was obtained. Those limiting values are summarized in Table I1 and Figure 2 (points marked -).
Table 11: Limiting Chemical Shifts of the Coordinated D M F Protons in D M F Solutions of AlXs and Calculated Proton-X - Distances --Cl--Signala
Au
a' b' C'
-Br--T
AV
T
b
...
12 35
2.96 2.28
5 10 30
3.72 3.12 2.38
--I---AU
4 8
18
T
3.93 3.30 2.70
Signal assignments correspond to those defined in Figure 1. Shifts, Av, are in Ha downfield from the corresponding signal from the complex A1(DMF)eaf. Bond distances, r, are in bngstroms. For high concentrations of LiC1, signal a' is shifted underneath signal b. a
Table I : Chemical Shifts (60 MHz) of D M F in 0.1 m Solutions of Aluminum(II1) Salts in D M F a t 37"
Salt
Al(ClO4)a AlCla AlBra AlIa
-------Chemical Free -methyl-a b
83 83 83
83
shifta------Coordinated --methyl-a' b'
93 93 93 93
102 103 104 105
112 118 119 118
Free formyl
Coordinated formyl
(0)
(0')
398 399 399 399
415 448 440 430
a Chemical shifts, in Hz ( k l Hz), downfield with respect to the internal standard cyclohexane. Assignments correspond t o the proton sites defined in Figure 1.
A study was made of the shift of the coordinated formyl proton resonance as a function of X- ion concentration at constant ionic strength using DMF solutions prepared from mixtures of Al(C104)3 and the appropriate AI&. A plot of the shift of the coordinated formyl proton resonance, relative to its position in Al(C10J3 solutions, vs. the [X-1: [Al(III)] ratio is shown in Figure 2. Data obtained a t other ionic strengths (0.2 5 I 5 1.2) are similar t o those summarized in Figure 2. The shifts obtained at equivalent [X-] : [Al(III)] ratios for solutions prepared from Aland LiX deviate upfield from the plots in Figure 2 by more than a simple ionic strength effect. However, by using those solutions, the [X-] : [Al(III)] ratio The Journal of Physical Chemistry
Unless one invokes a coordination number larger than 6 for Al(III), the halide ions in these systems must occupy positions outside the primary coordination sphere. Using the Buckingham modelL2for the chemical shifts of protons in an electric field (in this case caused by the X- ions), we have calculated the hydrogen-X- distances (see Table 11)from the limiting values of the chemical shifts. I n the calculation it was assumed that all coordinated DMF molecules are affected equally and that only the effect of the nearest halide ion need be considered (which seems reasonable in view of the l/r4 dependence of the shift).13 This model neglects any contribution t o the shift due to the magnetic anisotropy induced in X- by Al(II1) ; however, anisotropic effects should give rise t o an upfield contribution to the shift and result in even shorter (11) W. G. Movius and N. A. Matwiyoff, Inorg. Chem., 6, 847 (1967). (12) A. D. Buckingham, Can. J. Chem., 38, 300 (1960),and references therein. (13) Unfortunately, the stoichiometry of the ion pair is not known (see text), However, the assumption that all coordinated DMF molecules are affected equally leads t o the maximum calculated values for the Hi-X - distances. For example, if the structure is such that only half of the coordinated DMF molecules at a given time feel the presence of X- in the ion pair (which is formed rapidly on the nmr time scale), then the calculated Hi-X- distances would be -15% less than those listed in Table 11.
3065
NOTES
Figure 2. Plot of chemical shifts of the Coordinated formyl proton 0, X i s C1-, I = 0.6; 0, X i s Br-, Z = 0.9; A, X i s I-, Z = 0.8.
calculated H-X- distances.14*16 Although the calculated bond lengths may have no quantitative signific a n ~ e , the ~ ~magnitude )~~ of and trends in the values suggest that X- resides in the second coordination sphere of Al(II1). I n particular, we note that the A1(DMF)63+ion is quite hindered sterically and that, if X- approaches an octahedral face or edge of the complex in its preferred conformation, the Hi-X- distance should decrease in the order H,t-X- > Hbt-X- > H,,-X-. Efforts were made to determine the equilibria involved in the ion-pair formation by computer fitting the [X-] dependence of the formyl proton shift (Figure 2) to mass-action expressions for the reactions A1(DMF)ea+
+ X- +{ [A1(DMF)e3+]X-)’+
AI(DMF)6*+ -1 2X- + { [AI(DMF)6”]2X-) A1(DMF)e3+
+
(1) (2)
+ 3X- +{ [ A ~ ( D ~ ! I F ) B ~ + ] ~(3) X-}
The curves in Figure 2 could not be fit satisfactorily for the case in which mass-action expressions for reactions 1, 2, or 3 were considered individually or for the case in which mass-action expressions for both reactions 1 and 2 were treated.I6 Other combinations of equilibria possible were not examined because they involve a fit using more than two parameters (the limiting shifts and the equilibrium quotients). The uniqueness of
VS.
the ratio [X-] : [Al(III)] at a constant ionic strength (I):
the fit obtained under these conditions would be questionable, particularly when it is noted that the ionic strengths of the solutions used are quite high, and, in the treatment of the data, it was necessary to assume that no { [A1(DMF)Ba+]C104-] ion pairs are formed. We have reported previously8 the ‘7AI chemical shift ion in D M F solutions of Al(C104)a. of the AI(DMF)B~+ This chemical shift is unaffected by X- ion, but the 27Alline width does increase slightly (5-10%) as the [X-] : [A1(DMF)s3+]ratio is increased. Although the changes in line width are too small to allow a quantitative treatment, they are consistent with the lowering of the electric-field symmetry at AI(II1) due to ion-pair formation. The results presented here are of interest in connec(14) The magnitude and sign of the anisotropic contribution depends on the orientation of the coordinated D M F with respect to X-. Because the Al(DMF)sa+ ion is quite hindered sterically, D M F can adopt only a few “locked” conformations. For these conformations, anisotropic effects from X - should be shielding a t the protons. A quantitative assessment is not possible because we do not know what magnetic anisotropy, i f any, should be ascribed to the Al(II1)X - “bond.” Upper limits to the effect at the formyl proton of -30 HZfor I - and -20 Ha for C1- can be calculated if we assume that the anisotropy is as large as that for covalent C-X bonds.16 Thus it appears that the anisotropic effect in this system should be small. (15) H.M.McConnell, J. Chern. Phys., 27, 226 (1957). (16) ,Details of these calculations will be available in the Ph.D. Thesis of W. G. M. Volume 76,Number 8
August 1968
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.
