H. M. NELSON
1380
“Ludox” sol (although much less than in many other hydrosols), the rather crude estimate of the surface potential. However, what seems to be significant and encouraging is the fact that the theoretically determined Bzovalues are anywhere close to those experimentally determined and that they follow the expected trend. The influence of the coefficient AIoon the particle size of “Ludox” is not large, since the third root of the factor (1 jNaoA?)2has to be applied in corThe largest difference, amounting to recting D,.
+
Vol. 66
2%, is obtained for CNaCl = 0.3 M . This falls into the limits of experimental error of particle size determination, and to verify the influence of virial coefficient AI0 more precise experimental data would be needed. The experimental part of this work has been done in the_Department of Applied Biochemistry, “Andrija Stampar” School of Public Health, Faculty of Medicine, University of Zagreb, Zagreb, Croatia, Yugoslavia.
THE INFRARED SPECTRUM OF METHYL CHLORIDE IK STANNIC CHLORIDE AKD ANTIMONY PENTACHLORIDE SOLUTION BY H. M. NELSON^ Mellon Institute, Pittsburgh is,Pennsylvania Received November 8 , 1961
The infrared spectrum of CHaCl in SnCll solution (at 30 and -40’) and in SbC15solution (at 28 and -12’) has been obtained in an effort to elucidate the structure of the CHsCl:SnC4 and CH&l:SbCL addition compounds. No evidence of compound formation was found in the CH3Cl-SnC4 solutions. A tem erature-dependent band a t 688 em.-‘ in the CHICl-SbC15 solutions is assigned to the C-C1 stretching motion in the CH&l:SbC16 addition compound. The spectra of the CHaC1-SbCla solutions are consistent with a linear C-ClSb bond in CHsCl: SbC4 but the evidence is insufficient to rule out an angular bond which seems more probable from other considerations. The results suggest that the CHaC1:SnClr and CHsCl: SbCls addition compounds are more accurately described as slightly polarized complexes than as ion pairs.
Introduction Brown, Eddy, and Wong have suggested2that the initial step in the Friedel-Crafts reaction involves the formation of an addition compound (which may or may not undergo ionization) between the alkyl halide and the metal halide catalyst.
and that the distinction between the two may be somewhat arbitrary. They also suggest, since the alkyl group’s ability to accommodate a positive charge increases as it becomes more branched, that a t some point in the series-methyl, ethyl, isopropyI, t-butyl-the ion pair description becomes more accurate. This discussion raises a question RX 11x0 I_ R X R4Xa as to the degree to which the alkyl halide halogen is transferred to the metal halide in the addition In an effort to clarify the role of the Lewis acid compounds, Le., are they better described as catalyst, Brown and his co-workers have obtained polarized complexes or as ion pairs? The vibravapor pressure-phase composition data for a number tional spectra should provide information releof alkyl halide-metal halide s y ~ t e m s . ~Of $ ~par- vant to this point. Taylor and Moyere have exticular interest to the present work is Byrne’s amined the Raman spectra of a number of Lewis finding4 that methyl chloride forms 1:1 compounds acid-base systems, including methyl bromide-aluwith stannic chloride (CH3C1: SnC14,dissoc. press. minum bromide. They assign a band a t 554 40.30 mm. at -64”, calcd. m.p. -50°, heat of for- cm.-1 to the C-Br stretching motion in the CH3mation -4.69 kcal./mole) and antimony penta- Br:AIBrz addition compound, while a band at 594 chloride (CH3C1:SbC15, dissoc. press. 6.50 mm. at cm.-l is assigned to the corresponding motion in -50°, calcd. m.p. 90°, heat of formation -8.92 the CH3Br present as solvent. The objectives of kcal./mole) . the present work were to obtain infrared spectra of Jungk, Smoot, and Brown5 present two alterna- some alkyl halide-metal halide addition compounds tives for the mechanism of the aluminum bromide- with particular attention t o the effects of compound catalyzed alkylation of benzene and toluene: an formation on the alkyl halide spectrum and to interattack by the aromatic on a polarized addition pret these effects in terms of the structure of the compound R+Br-:A1Br3 or reaction of the aroma- compound formed. tic with an ion pair intermediate, R+A1Br4-. They Experimental point out that there’appeared to be no experimental basis for a choice between the two formulations Materials.-The materials used and the methods for
+
(1) Chemical Products Division, Aerojet-General Corporation, Azusa, Calif. (2) H. C. Brown, L. P. Eddy, a n d R . Wcng, J . Am. Chem. SOC.,75, G275 (1953). ( 3 ) (a) H. C. Brown and W. J. Wallace, ibid., 75, 6279 (1953); (b) R. TVong and H. C. Brown, J. Inorg. Nuel. Chenz., 1,402 (1955). (4) J. J. Byrne, Ph.D. Thesis, Purdue University, 1958; Dzssertation Abstr,, 18, 1976 (1958). ( 5 ) H. Jungk, C. R. Smoct, and H. C. Brown, J. Am. Chem. Soc., 78, 2185 (1956).
their purification and subsequent handling followed those of the work already m e n t i ~ n e d . ~ ~ ~ Infrared Cells.-The moisture-sensitivity and generally reactive nature of the materials involved required modification of conventional infrared sampling techniques. The celIs (6) R. C. Taylor and J. R. Moyer, Abstracts, 135th Meeting of American Chemical Society, Boston, April 1959; J. R. Moyer, Ph.D. Thesis, University of iMiohigan, 1958, Dissertation Abstr., 19, 3149 (1959).
August, 11362
INFRARED SPECTRUM OF METHYL CHLORIDE IS STANXIC CHLORIDE
1381
used consisted of a pair of potassium bromide windows TABLE I separated by a IJ-shaped Teflon (Du Pont) spacer. The CHIC1 ABSORPTION FREQUEKCIES (cm. -9 windows were held together by a strip of epoxy cement SnClr SbCls (Hysol Epoui-Patch Kit 1C) applied over the line joining Vapora soh. soln. them. (Over-all dimensions of the cell were 16 X 48 X 10 nim. with a working cell area of 6 X 43 mm. Cell thicknesses v , ( ~ str., 714 ~ A,, ~ ) 732.1 716 ranged from, 0.055 to 0.250 mm.) The open end of the cell 688' proper then was butt-jointed with a fillet of the same cement I018 1015.0 1014 to a Pyrex U-tube (volume approx. 35 ml.) fitted with side- vg(CH3 rock, E, ,I) arms. Theee cells were easy to construct and proved to be w(CH3 def., AI, /I ) 1343 1354.9 1346 vacuum tighk. 1434 H S E, I ) 1454.6 1441 Sample Introduction.-In order to introduce the sample, ~ ~ ( C def., 2848 2878 2850 the U-tube was attached to a separate branch of the vacuum 2 V d 2953 2966 2955 manifold through a sidearm. It was separated from stop- vdCH str., AI, cock grease and mercury by a liquid nitrogen-cooled trap. v4(CH str., E, I ) 3027 3041.8 3026 A section of tubing to which samples of the metal halide and a Reference 7, p. 313. I n CW&X:SbClb. methyl chloride were attached (on either side of a second Utube trap) was sealed to the other end of the cell U-tube. After evacuation, the fragile tip on the metal halide sample Lowering the temperature to -40' (Fig. 1) was broken and the metal halide was distilled into the liquid produces a general increase in intensity, presumably nitrogen-cooled cell I!-tube. Then the trap between the methyl chloride sample and the cell U-tube was cooled to as a result of the increased solubility of the methyl -80" and the methyl chloride sample WM introduced. This chloride. There is also a slight change in the relatrap rJerved to eliminate traces of mercury contained in the tive intensities of v 2 and v s (the CIl3 deformation), methyl chloride sample. The section of tubing containing No new features are evident. In other samples in the empty sample ampoules then was sealed off and the cell U-tube itself was sealed off from the manifold. After the which v3 (the C-C1 stretch) was not totally absorbsample had warmed to room temperature, the apparatus was ing a t the lower temperature, no splitting was obinverted to allow the sample to run from the cell U-tube into served. the infrared cell proper. A reference cell of the same thickThe spectrum of methyl chloride dissolved in ness was filled with the metal halide solvent in a similar manner. The concentrations of the methyl chloride solu- antimony pentachloride a t 28' is shown in Fig. 2 . tions were unknown because of the unknown distribution of The general appearance of the spectrum is the same the two components between the liquid 2nd vapor phases. as that in stannic chloride, all of the bands being The quantities of materials used were on the order of 1mmole shifted downward from their vapor state values of the, methyl chloride and 1-15 mmoles of the metal halide. The most useful spectra were obtained with mole ratios of with the exception of the methyl rocking band. about 1:l and 0.120-mm. cells. The cells resisted attack In addition a new band appears a t 688 em.-' by thesample to different degrees. With the SbC16 solutions, while v3 a t 714 em.-l is relatively weaker than in discoloration wm evident after a few days and unidentified SnC14 solution. On going to lower temperatures weak new bands appeared in the spectrum. The 8nC14 solutions seemed t o be stable for much longer periods. With (Fig. 2 ) , the band a t 688 cm.-l appears to grow in both kinds of Solutions, bands, presumably due to water, intensity a t the expense of the 714-cm.-l band. began to appear a t about 1580 and 3450 cm.-I after several There is a relative change in the intensities of v2 weeks. Thiai type of cell would seem to be generally useful in obtaining spectra of volatile, moisture-, or ouygen-sensi- and vs which is more pronounced than that in SnC14 solution. There is no evidence of splitting tive materials. Spectra.-Spectra were obtained with a Eeckmnn IR-4 of V 6 , the methyl rocking band, nor of any of the instrument using sodium chloride and cesium bromide optics. other doubly degenerate vibrations. The temperaThe absorption frequencies were checked on a Perkin-Elmer ture dependence of the 688-cm. band suggests 112 instrument equipped with sodium chloride and calcium fluoride prisms. The accuracy of the frequencies is estimrted that it is associated with the CH3C1:SbCls addition compound. Lowering the temperature would be to be & l cm +for thoae below 2000 cm.-l and A 5 cm.-'for those above. Spectra were obtained from 400-3500 cm.+. expected to shift the equilibrium No absorptions were observed in either SnC1, or SbCls sohtion in the I-egion 400-600 cm.-'. The low temperature CH3C1 4-SbCIs CH3Cl:SbCIs spectra were obtained by placing the cell in an enclosure constructed of blocks of Styrofoam (Dow) cemented together. in the direction of the addition compound. BeThis enclomre was slotted to allow passage of the infrared beam. Two "manifolds" were cut into the base of the en- cause of this and in view of its position, the 688closuria with a cork borer in such iz way as to allow one stream cm.-' band is assigned to the C-C1 stretching in of nitrogen (Linde H.P. Dry, cooled by paissage through a CHsCl :SbCls. trap immersed in Dry Ice or liquid nitrogen) to flow over the Jones and Sheppard8 studied the spectra of faces of the cell, while a second stream (at room temperature) inore removed from the cell faces prevented condensa- methyl bromide and methyl iodide in carbon tetration or entrainment of atmospheric moisture. The cell chloride solution. They found that the parallel temperature was measured with a thern~ocoupleinserted in a bands VI and v% (v3 was not observed) showed hole drilled near the edge of one of the cell windows. This comparatively small half-widths while the doubly device permitted cooling the cell to ternpcratures as low as -40" for periods of 0.5 hr. or more with no trouble from degenerate perpendicular bands v4. vs and V6 fogging of the cell faces. The temperature of the cell could were considerably broader with extensive wings. be kept constant to within a few degrees by controlling the They attribute this difference in behavior to a depth of immersion of the trap through which the nitrogen restriction of the rotation about axes perpendicular passed for ccding.
11
Ill
Results and Discussion The spectrum of methyl chloride dissolved in stannic chloride a t 30" is shown in Fig. 1. The bands are shifted downward 10-30 cm.-l from the values in the vapor state with the exception of YE (the CH3rocking), as shown in Table I.
I/)
to the symmetry axis (the carbon-halogen bond) while rotation about this axis remains free or quasi-free. The same general behavior is apparent in the spectra discussed here. Of particular (7) G. Hersberg, "Infrared and Raman Spectra," D. Van Nostrand Co., Inc., New York, N. Y.,1945. (8) W. J. Jones a n d N. Sheppard, ?'Tuna. Faraday Sac., 56, 625 (1960).
1382
H. M. NELSOX
Vol. 66 I
100,
_I
3500
I
'
3000
1
;
2500
2OW
I
I
I
I
I
I
I
I
1
1
1
I
1500
FREQUENCY IN
Fig. 2.-Infrared
1
I
1
I
1000
I
Ill
I
0
CW1.
spectrum of methyl chloride in antimony pentachloride solution: 28" -,
interest in this connection are the 688 and 714 cm.-l hands in SbClb solution. While they are not well resolved, the former, which must be essentially a vibrational motion, appears to have about the same half-width as the latter, thus supporting Jones and Sheppard's views. The intensity of the 688-cm.-l band in the CH3Cl-SbC16 solutions indicates that a significant portion of the methyl chloride present is in the form of the addition compound CH3C1:SbCla. I n spite of this, no splitting of the doubly degenerate modes is observed. It might be argued on this basis that the methyl chloride effectively retains the CBVsymmetry of the free molecule in the addition compound, that is, that the C-C1-Sb bond is linear. If this bond were bent, the symmetry would be expected to be reduced to C. with a con-
I
- 12' - . -.
sequent splitting of the doubly degenerate E species bands. On the other hand, the splitting may be so small as to escape detection, especially in the presence of dissolved but "free" methyl chloride. Thus while the observed spectrum is consistent with a linear structure, the evidence is insufficient to eliminate an angular one. KO plausible disposition of the chlorine's bonding electrons gives rise to a linear C-C1-Sb bond. If the chlorine uses four u-electron pairs in sp3 hybrid orbitals, mith two lone electron pairs (as in ClOz-), the resultant complex would be V-shaped with a bent C-C1-Sb This alternative seems more attractive, even in the absence of evidence from the spectra. If the spectra are considered in terms of a "tri(9)
R. J. Gillespie and R. S. Nyholm, Quart. Re%.,ll, 373 (1957).
August, 1962
ION-EXCHANGE RESIXABSORPTION O F BROMIDE ION
atomic” RXRl model for the addition compounds (where El is the methyl group, X the chlorine, and M the l3bClb or SnCL molecule) and a simple valLenceforce field,10 it is found that the effect of addition compound formation (regarded as the establishment of an XM bond) on the RX stretching frequency is strongly dependent on the change in the RX stretching force coiistant on compound formation. The calculation of the RX force constant in terms of this model would require knowledge of the XM stretching and RXM bending frequencies. No attempt was made to observe these bands in view of their probable lorn frequencies. For the same reason, the effects of compound formation on the metal halide spectra were noit examined. The possibility of an ion pair structure, CHs+SblCl6-, for the SbC16 addition compound seems unlikely in the absence of any pronounced changes in the methyl group frequencies, jn spite of the fact that a large proportion of the CH3C1present is in the form of the addition compound. I n addition, the decrease in intensity of the 714-cm.-l band which occurs on lowering the temperature (IO) Reference 7, p. 173.
FROM
SALT
SOLUTIONS
1383
appears to be balanced by the increase in the 688cm.-l band, which would be improbable if an ion pair structure also mere participating in the equilibrium. The failure to observe any clear evidence of compound formation in the SnCL solutions may be due to the instability of the CH3C1:SnCld addition compound itself, or to a coincidence of the C-Cl stretching frequency in the addition compound vcvith that in the free, dissolved molecule. The possibility of the formation of an ionic compound seems unlikely in view of the results in SbCls solution and the relative instability of the SnC14 addition compound. The experimental results suggest that the CH3C1: SbC16 and CH3C1: SnCle addition compounds are more accurately described as slightly polarized complexes than as ion pairs. Acknowledgments.-This work was made possible by a grant from the U. S. Army Research Office, Durham. This support is gratefully acknowledged. The author also is indebled to his colleagues, Drs. D. E. Milligan and IT. G. Fateley, for helpful discussions.
LIQUID AMINE ,4ND IOS-EXCHANGE RESIX ABSORPTION OF BROMIDE ION FROM AQUEOUS SALT SOLUTIOSS; VARIATION WITH AQUEOUS ELECTROLYTE ACTIVITIES1 BY S. LISDENBAUM AND G. E. BOYD Chemistry Diviszon, Oak Radge National Laboratory, Oak Rzdge, Tennessee Receited November 8. 1961
The absorption of microquantities of bromide ion from aqueous solutions of hydrochloric acid, lithium, sodium, potassium, and cesium chlorides by tri-n-octylamine solutions in toluene, and from hydrochloric acid by a strong-base anion exchanger was measured. The departure of the equilibrium distribution coefficients for the liquid ion exchanger from a “mass law” dependence on electrolyte concentration could be explained quantitatively by the non-ideality of the aqueous phase except for the concentrated hydrochloric acid solutions. With the anion-exchange resin, however, corrections for aqueous phase non-ideality and for de-swelling were not sufficient to cause the distribution data to conform to the mass law; in addition, an important correction for the organic hase non-ideality was required. The magnitude of the organic phase non-ideality could be correlated with the invasion or this phase by electrolyte. \Then the invasion was small, cr absent, as with the liquid exclianger, the non-ideality was small and could be neglected.
The role of the secondary cation on the absorption of anionic species from concentrated aqueous electrolyte solutions by strong-base anion-exchsnge resi.ns has been the subject of several studies.2-6 It was observed by Kraus and co-workers6J that many m.et,al chloro-complex anions were more strongly absorbed from LiCl than from HC1 solutions by Dowex-1. This phenomenon was shown by Horne2 to be of more general occurrence. This worker and others found, however, that the concentration variation of the distribution coef(1) This paper is based upon work performed a t Oak Ridge National Laboratory, which is operated by Union Carbide Corporation for the Atomic Energy Commission. (2) R. A. Horne, J . Phys. Chem., 61,1651 (1967). (3) B. Chu a n d R. M. Diamond, {bid.,63, 2021 (1959). (4:l Y. Marcus a n d C . D. Coryell, Bull. Res. Council Isvael. A8, 1 (1959). (5) K. A. Kraus a n d F. Nelson, Proe. .Tntern. Conf. Peaceful Uses A t . Energy, 7,113 (1956). (6:i K. A. Kraus, F. Nelson, F. B. Ciough, and R. C. Carlston, J . A m . Chem. Soe., 7 7 , 1391 (1955).
ficients for the ZnC14-2 anion in acid and alkali metal chloride solutions did not give slopes of -2 in log-log plots as expected from the application of the simple law of mass action. It has been suggested3+ that the non-ideal variation of the distribution coefficients for this and other anions with the nature and concentration of the aqueous electrolyte was caused by special effects in the organic phase. Schindewolf7 compared the extraction of Zn(I1) from aqueous HC1, LiC1, and CsCl solutions by a quaternary ammonium type anion-exchange resin with that by methyldioctylamine dissolved in trichloroethylene, and noted that the variations of the distribution coefficients with the aqueous electrolyte activity were strikingly similar. The fact that the departures from ideality appeared to be the same for the liquid and resin ion-exchange systems led Schindewolf t o suggest that the variation (7) U. Schindewolf, 2. Elelctrochem., 6 2 , 336 (1958).
