RYOHEINAKANE AND TOSHIYUKI OYAMA
1146
Acknowledgments. The authors gratefully acknowledge the assistance of Charles F. Jumper and William J. Sutton, who did the preliminary work on this prob-
lem. The project was partially supported by the Atomic Energy Commission under Contract No. AT(40-1)-1437.
Boron Isotope Exchange between Boron Fluoride and Its Alkyl Halide Complexes, 11.1 Infrared Spectrum of Boron FluorideMethyl Fluoride Complex
by Ryohei Nakane and Toshiyuki Oyama The Institute of Physical and Chemical Research, Bunkyo-ku, Tokyo, Japan
(Received October 15, 1965’)
The infrared spectra of liquid B10F3,B1*F3,CH3F, B10F3.CH3F complex, and Bl1F3.CH3F complex are observed in the region from 400 to 4000 cm-l. The isotopic data are used to calculate the theoretical equilibrium constant for boron isotope exchange between gaseous boron fluoride and its methyl fluoride complex. The calculated values are found to agree fairly well with observed values.
Previously, one of the authors found that, for boron isotope exchange, the equilibrium constant is much smaller between gaseous boron fluoride and weak boron fluoride complexes in the liquid form than between gaseous boron fluoride and strong boron fluoride complexes in the liquid form. Namely, with boron fluoride-alkyl halide complexes1 or boron fluoride-alkyl halide-alkylbenzene 1: 1: 1 addition oriented 7~ complexes,2 which exist only a t low temperatures, the constant is much smaller than with boron fluoride-ether complexes which are stable even at room temperature. The known infrared spectra of gaseous boron fluoridea and liquid boron fluoride-ether complexes4 were used to calculate the theoretical equilibrium constant for boron isotope exchange. I n the present work, the infrared spectrum of liquid boron fluoride-methyl fluoride complex was observed a t low temperatures and the equilibrium constants, observed and calculated from isotopic data of infrared spectrum, were compared. The Journal of Physical Chemistry
Experimental Section For the measurements of infrared spectra at low temperatures a cell as shown in Figure 1 was made. The specimen to be studied was introduced into the space between two KRS-5 plates (1 cm diameter by 5 mm thick) A and A’, between which spacer B of Teflon was inserted to give a space of about 0.05 mm thickness. The t,wo KRS-5 plates were firmly held in brass holder C with two Teflon gaskets D and D‘ for vacuum tightness. The cryostat was a stainless steel dewar with two KRS-6 plate windows (2 cm diameter by 3 mm thick), E and E’. The thickness, (1) Part I : R. Nakane, 0. Kurihara, and -4. Natsubori, J . Phys. Chem., 68, 2876 (1964). (2) R. Nakane, A. Natsubori, and 0. Kurihara, J . Am. Chem. Soc., 87, 3597 (1965). (3) J. Vanderryn, J . Chem. Phys., 30, 331 (1959). (4) A. A. Palko, G. M. Begun, and L. Landau, ibid., 37, 552 (1962); G. M. Begun and A. A. Palko, ibid., 38, 2112 (1963); G.M. Begun, W. H. Fletcher, and A. A. Palko, Spectrochim. Acta, IS, 655 (1962).
INFRARED SPECTRUM OF BORON TRIFLUORIDE-METHYL FLUORIDE COMPLEX
to gas reservoir through vacuum cock
to
_ _ _ _
vacuum
1147
pentane at its melting point (- 130.8') was poured into tube F, the solid specimen in C melted, and at this instant the thermocouple G attached on the outer surface of C indicated approximately -105'. The cell window was placed directly in the radiation path of a Perkin-Elmer RiIodel 521 grating infrared spectrometer. Enriched B1"F3was obtained by multiplying boron isotope exchange between boron fluoride gas and boron fluoride-anisole complex ; it was analyzed as 95 atom % B10.5 In place of enriched B11F3, normal boron fluoride was used. Jlethyl fluoride was prepared from methyl tosylate by reaction with potassium fluoride.6
Results and Discussion The infrared spectra of liquid B1°F3, B11F3, CH3F, B1"F3.CH3F complex, and B"F3 CH3F complex were observed in the region from 400 to 4000 cm-l. Our infrared spectrum of liquid boron fluoride agreed fairly well with the known infrared spectrum of solid boron fluoride' as shown in Table I and Figure 2. Hence the 1416-cm-l frequency in liquid B11F3 can be assigned to B-F antisymmetric stretch. I n liquid B11F3.CH3F complex the absorption band was observed a t 1424 cni-1 as shown in Table I1 and Figure 3. I n the region from 1000 to 1424 cm-', no frequency other than the frequency of combination (vl v4) in boron fluoride and of CH3 rock was observed. On the other hand, the B-F antisymmetric frequencies in BllF3. CH30CH3complex3 and B11F4- ion* were observed a t 1177 and 1216 em-' for the former and near 1050 cm-1 for the latter. Therefore, the 1424-em-' frequency in BI1F3.CH3Fcomplex is assigned to the B-F antisymmetric stretch and no BF4- ion is contained in liquid BF3.CH3F complex. The change in frequency with complex formation was not observed in the BF3 in-plane bending band. The out-of-plane bending band in liquid B"Fs was observed at 661 cm-l, but in liquid B11F3.CH3F complex no band was observed there, and a new band was observed a t 617 em-' which can be assigned to out-of-plane bending band in liquid Bl1F3.CH3F complex. The intensity of the band in the complex is higher than that in liquid BF,. The B-F symmetric stretching frequency, which is not infrared active, was not observed.
+
E
D' Figure 1. Low-temperature absorption cell.
3 and 5 mm, of KRS plates, though detrimental to infrared work, was necessary to assure the strength of the plates. Experiments were made as follows. After evacuating the cryostat to dewar vacuum, liquid nitrogen was poured into center tube F of the cell, the outer surface of which was coated with charcoal to hold dewar vacuum for a long time. The specimen gases were then introduced one after another through a vacuum cock into the cell and solidified in holder C. When n-
(5) R. Nakane and T. Watanabe, D6itai To Hdshasen, 2 , 273 (1959). (6) W. F. Edge11 and L. Parts, J . A m . Chem. Soc., 77, 4899 (1955). (7) D. A. Dows, J . Chem. Phys., 31, 1637 (1959). (8) (a) J. Goubeau and U ' . Bues, Z . Anorg. Allgem. Chem., 268, 221 (1952); N. N. Greenwood, J . Chem. SOC.,3811 (1959); G. L. Cot6 and H. W. Thompson, Proc. Roy. SOC.(London), A210, 217 (1951); (b) B. P. Suss and J. J. Wuhrmann, Hela. Chim. Acta, 40, 722 (1957) ; (c) D. Cook, S. J. Kuhn, and G. A. Olah, J. Chem. Phys., 33, 1669 (1960).
Volume 70,Number 4
April 1966
RYOHEINAKANEAND TOSHIYUKI OYAMA
1148
100
. se f
-
80
I
60-
2
-
'BI
40
i;
200, 4000
3500
3000
2500
2000
1800
1600
1400
1200
1000
800
600
400
Frequency, cm-1.
Figure 2. Infrared spectrum of liquid boron fluoride:
4MH)
3500
3000
-,
2500
B10F3;
, B"F8.
_ I _
1800 1600 Frequency, om-1.
2000
Figure 3. Infrared spectrum of liquid boron fluoride-methyl fluoride complex: ----, B11F3.CH3F complex.
1400
1200
1000
-, B1aF3.CH,F
800
600
400
complex;
100
80 iy
60
d
'3
40
20 0 4000
3500
3000
2500
2000
1800
1600
1400
1200
1000
800
600
400
Frequency, cm-1.
Figure 4. Infrared spectrum of liquid methyl fluoride.
The infrared spectrum of liquid methyl fluoride agreed fairly well also with the known spectrum of gaseous methyl fluorideg as shown in Table I11 and Figure 4. Thus, the bands observed at 1013, 2966, and 3018 cm-l in liquid methyl fluoride are assigned to C-F, C-H symmetric, and C-H antisymmetric stretches, respectively. I n both liquid isotopic complexes, these bands were observed a t 990,2980, and 3046 cm-1, which shows that the C-F force constant becomes slightly lower and the C-H force constant beThe Journal of P h y s i d Chemistry
comes slightly higher when methyl fluoride forms the complex with boron fluoride; there was no boron isotope effect on these force constants. The change in frequencies of the CH, bend and the CH3rock in methyl fluoride with complex formation was not distinct. The values of the equilibrium constant for boron isotope exchange were calculated from infrared spectra data, obtained on boron fluoride-methyl fluoride com(9) E.F.Barker and
E.K.Plyler, J . Chem. Phys., 3 , 367 (1935).
INFRARED SPECTRUM OF BORON TRIFLUORIDE-METHYL FLUORIDE COMPLEX
~~~~~
~
1149
~
Table I : Vibrational Frequencies of Gaseous, Liquid, and Solid BFa (cm-l) -Gaseous B 10
888 718.23 1504.8 482.0 See ref 3.
BFPBI1
-
888 691.45 1453.9 480.4
' This work.
-Liquid B 10
BFab-
-Solid B11
...
...
685 1467 489
661 1416 487.5
B'QFa*CHaF
B"Fa.CHsF
Aesignment
489 638 990 1182 1351
487.5 617 990 1182 1350
BF3 in-plane bend BF3 out-of-plane bend C-F stretch CH3 rock Combination band ( V I v4 of BFs) B-F stretch (antisym) (CHI bend) ( CH3 bend) CHI bend (overtone) C-H stretch (sym) C-H stretch (antisym)
... ... 2863 2980 3046
1424 (1451) (1483) 2863 2980 3046
Liquid CHaFb
1048.2 1195.5 1475.3 1471.1 2861.6 2964.5 2982.2
1013 1182 1464
'See ref 9.
2860 2966 3018
876-880 658 1457 474-484
876-880 632 1405 472-483
Assignment
vl, B-F stretch (sym) VZ,
YS
v4
BF3 out-of-plane bend (2), B-F stretch (antisym) (2), BF3 in-plane bend
the density of localized electrons on the vacant orbital of boron atom and hence the polarity of complex becomes higher, causing the equilibrium constant to become larger. Table IV : Equilibrium Constants for Boron Isotope Exchange between Gaseous BF3 and Liquid BF3.CHsF Complex
+
Table I11: Vibrational Frequencies of Gaseous and Liquid CHaF (cm-1) Gaseous CHaF'
BI1
See ref 7.
Table I1 : Vibrational Frequencies and Band Assignments for Liquid BFa.CH3F Complex (cm-1)
1475
BFacBlO
Assignment
C-F stretch CHa rock CHs bend (sym) CH3 bend (antisym) CHI bend (overtone) C-H stretch (sym) C-H stretch (antisym)
' This work.
plex and gaseous boron fluoride, by the use of the formulas given by Ureylo and Bigeleisen and Mayer, 11 assuming that there is no boron isotope shift of the B-F symmetric stretching frequency in the complex as in gaseous boron fluoride. It was found that the equilibrium constant depends mostly on the boron isotope shift of the BF3out-of-plane bending frequency. The calculated values of the constant were in fair agreement with the observed as shown in Table IV. Thus, the BF3 out-of-plane bending frequency in a weak boron fluoride complex will become lower as
P E q u i l constant, Calcd
Temp, OC
- 95 -112 a
K Obsda
1.020 -1.026
1.020 1.023
See ref 1.
I n a previous work,' one of the authors assumed that the B-F antisymmetric stretching frequency in boron fluoride decreases and the structure of boron fluoride deforms slightly from planar towards tetrahedral when boron fluoride forms the weak complexes with alkyl halides and that this change results in the increase of values of the equilibrium constant for boron isotope exchange. However, the above assumption was found mistaken. When liquid boron fluoride is dissolved in liquid methyl fluoride, a lone pair of electrons of the fluorine atom in the methyl fluoride is partially localized on the vacant orbital of the boron atom, causing (a) the force constant of BF3 out-ofplane bend to become lower, (b) the B-F force constant to become slightly higher, (c) the C-F force constant to become slightly lower, and (d) the C-H force constant to become slightly higher. The B-F antisymmetric stretching frequency, which degenerates doubly in the uncoordinated boron fluoride molecule, is split into two by the destruction of the molecular symmetry in the BF3.CHIOCH3 complex, but not in the BF3. CH3F complex. Hence, in the latter weak complex the weakening of the B-F bond does not occur and the ~~~~
~
~
(10) H.C. Urey, J. Chem. Soc., 562 (1947). (11) J. Bigeleisen and M. G . Mayer, J. Chem. Phya., 15, 261 (1947).
Volume 70, Number 4 April 1966
1150
PAOLO BELTRAME, S. C A R R AND ~ , S. MORI
boron fluoride retains the planar structure. The increase of the values of the equilibrium constant for boron isotope exchange in the weak boron fluoride complex system results from the lowering, not of the B-F force constant, but of the force constant of the BF, out-of-plane bend. As for the weakening- of the C-F bond, it is very small and the methyl carbon retains a tetrahedral geometry. Thus, the ionic carbonium
form is absent in the weak boron fluoride complex as in CH,CI.SbCI, complex.'*
Acknowledgment. The authors wish to thank Mr. Teruo Kurihara for helpful advice on the infrared analyses. (12) H. M. Nelson, J. Phys. Chem., 66,1380(1962).
Competitive-Consecutive Reactions in the Photochemical Chlorination of p-Xylene
by Paolo Beltrame, Sergio Car& and Sandro Mori Istituto d i Chimica fisica, Universe'tb di Mitano, Italy
(Received October 18, 1966)
A kinetic study of the side-chain chlorination of pxylene was carried out in CCL solution at 30 and 50" under ultraviolet irradiation. Seven products from monochloroto hexachloro-p-xylene were detected and determined by gas chromatography. On the basis of a system of kinetic equations first order with respect both to chlorine and to the organic compounds, relative rate constants, referred to the specific rate of p-xylene chlorination, were evaluated. A solution of the set of kinetic equations was obtained such as to give reagent and products concentrations as functions of the chlorination degree. Employing the values of the relative rate constants (at 50'), a good fit was obtained of calculated curves of concentrations vs. chlorination level and observed values. Taking into account statistical factors, relative rate factors, fr, for attack of chlorine atoms to single C-H bonds were derived. The fr values prove that mainly the negative inductive effect of chlorine substituents governs the rates of the reactions. A comparison is made of the relative rates constants a t 30 and 50'.
The formal kinetics of complex reaction systems has been the object of extensive mathematical research;lJ! however, the proposed calculation schemes have been applied to relatively few cases, for lack of experimental data. A suitable reaction system is the side-chain chlorination of alkylbenzenei. The simple case of toluene has been by particularly by Haring and Knol.a A xylene, for instance, the para The Journal of Phy8e'Cal C h m k t r y
isomer that we have chosen, can on principle give the pattern of competitiveconsecutive reactions shown in Scheme I. (1) J. Wei and C. D. Prater, Advan. C ' a t d y s ~ ,13, 203 (1962), and references cited therein. (2) N. M. R o d i d n and E. N. Rodiguina, "Consecutive Chemical Reactions," D. Van Nostrand Co., Inc., New York, N. Y.,1964. (3) H. G. Haring and H. W.Knol, Chem. Process E w . , 45,560,619, 690 (1964); 4 6 , 3 8 (1966).
