R. NAKANE, 0. KURIHARA, AKD A. SATSUBORI
2876
Boron Isotope Exchange between Boron Fluoride and Its Alkyl Halide Complexes. I.
Relation between Equilibrium Constant of Isotopic
Exchange Reaction, Stability, and Catalytic Activity of Boron Fluoride Complex
by Ryohie Nakane, Osamu Kurihara, and Akiko Natsubori The Institute of Physical and Chemical Research, Bunkyo-ku, Tokyo, Japan
(Received December SO, 1963)
Equilibrium constants of boron isotope exchange between boron fluoride gas on one side and boron fluoride-t-butyl chloride, -isopropyl chloride, and -methyl fluoride complexes on the other are found as 1.033, 1.021, and 1.026 a t - 112'. Toluene is found to be readily alkylated with t-butyl chlorlde as well as alkyl fluoride in the presence of boron fluoride. Moreover, it is suggested that if the equilibrium constant becomes large, the polarity of the complex will increase; hence, the stability and catalytic activity of the complex will become high. This leads to the following conclusion : boron fluoride-t-butyl chloride, -methyl fluoride, -methyl chloride, and -isopropyl chloride complexes are polar complexes, 6-
6+
and the polarity of these complexes decreases in the order: BFa+ClC(CHa)3 6-
6-
6+
6+
6-
>
6+
BF3+FCH3 > BF3+ClCH(CH3)tl BF3+C1CH3. The aromatic constituent reacts with boron fluoride- t-butyl chloride complex by a nucleophilic displacement reaction.
Introduction Aromatic compounds are alkylated with alkyl halides in the presence of metal halides (Friedel-Crafts catalysts). Brown, et u Z . , ~ proved that in these reactions the aromatic constituent reacts with a polar alkyl a+
6-
halidemetal halide complex, RX-+MX3 by a nucleophilic displacement reaction, although for tertiary alkyl halides which are easily ionized, they assumed the formation and reaction of a different carbonium ion complex, R+MX4-. The existence of metal halide-alkyl halide complex& has been proved by many studies. Olah, et aLI2found polar alkyl fluoride-boron fluoride complexes by measuring the specific conductivity a t low temperatures. These complexes can alkylate readily aromatic conipounds, but it has been believedYthat aromatic compounds cannot be alkylated with alkyl chloride in the presence of boron fluoride, the cause of which has been explained as due to the greater stability of BFI- than BF3C1-. Recently, some of the authors proved by measuring the equilibrium constant of isotopic exchange and the The Journal of Physical Chemistry
absorption apectrum that boron fluoride can form acomplexes only with polar monoolefins such as 'propylene14l-butene,s 'and cis-2-butenes a t low temperatures, while the a-complex formation with nonpolar monoolefins such as ethylene4 and trans-2-butene5 cannot be found except only as very unstable a-complexes. Some of the authors found also that the equilibrium constant of isotopic exchange between boron fluoride gas .and the boron fluoride-methyl fluoride complex is 1.020 'at - 9 5 O and believed that the exchange distillation of boron fluoride-methyl fluoride complex a t low temperatures is a most promising method for the separation of boron isotopes.6 ~~
(1) H. C.Brown, H. W. Persall, L.P. Eddy, W. J. Wallace, M. Grayson, and K. L. Nelson, Ind. Eng. Chem., 45, 1462 (1953). (2) G. Oldh, S. Kuhn, and J. Oldh, J. Chem. Soc., 2174 (1957); G.A. OlLh and S. J. Kuhn, J . A m . Chem. SOC.,80, 6541 (1958). (3) R.L.Burwell and S. Archer, ibid., 64, 1032 (1942); G.F. Hennion and R. A. Kurs, ibid., 65, 1001 (1943); G. A. Russell, ibid., 81, 4834 (1959). (4) R. Nakane, T. Watanabe, and 0. Kurihara, Bull. Chem. SOC. Japan, 25, 1747 (1962); R. Nakane, T. Watanabe, 0 . Kurihara, and T.Oyama, ibid., 36, 1376 (1963). (5) R. Nakane, T.Watanabe, and T.Oyarna, ibid., 37, 381 (1964).
BORONISOTOPE EXCHANGE BETWEEX BF8AND ITSALKYLHALIDECOMPLEXES
I n the present work, measurements of the equilibrium constants of isotopic exchange between boron fluoride gas and several boron fluoride-alkyl halide complexes were made and the alkylation of toluene in the presence of boron fluoride was examined with t-butyl chloride. The obtained equilibrium constants and catalytic activities of these complexes are compared with those of many other boron fluoride complexes, and the relation between: the equilibrium constants and the catalytic activities of boron fluoride complexes is discussed again in detail.
Experimental Materials. Commercial tank boron fluoride was purified by low temperature distillation in a column (40 cm. in length and 1.5 cm. in internal diameter) packed with 1.5-mm. Dixon gauze rings of stainless steel. Isopropyl chloride and t-butyl chloride were commercial products ; they were fractionated by repeated bulb-tobulb distillation i n vacuo. Methyl fluoride was prepared from methyl tosylates by reaction with potassium fluoride by the method of Edgell, et al.,7and purified by low temperature distillation in the column mentioned above. Toluene, Fhich was the purest commercial product obtainable, was used directly. Measurement of Isotope Eflect. Experimental methods were exactly the same as those described in the previous papers. Determination of single-stage separation factor is made by the following procedure. Known amounts of purified boron fluoride and alkyl halide were introduced into a flask of approximately 300 ml. with a break-seal. The flask was closed and immersed in a low temperature bath at the melting point of carbon disulfide ( - 112') or toluene (-95'). A two-phase system composed of liquid boron fluoride-alkyl halide complex and gaseous boron fluoride was then produced in the flask. From the known volume of the flask and the total amounts of introduced boron fluoride and alkyl halides, m moles of boron fluoride in the gas phase and M moles of boron fluoride in the liquid phase were calculated. Since a large excess of boron fluoride jn liquid phase was used, the determination of single-stage separation factor was made only by measuring the BIO/B1l ratio in the gas phase before and after the equilibration. The flask was vigorously shaken by a shaker for approximately 2 hr., after which boron fluoride gas in the flask was sampled and directly analyzed with a mass spectrometer, and the isotopic ratios of B1"to BI1 at the peaks for mass 10 and mass 11 were measured. From the analyzed values corrected by the effect of m / M , the single-stage separation factor 01 was obtained. The error in 01 was fO.001, since a precision of about 4vy
2877
fO.0005 for the isotopic ratios of BlUto Bll was obtained from mass analyses repeated 40 times. Alkylation of Toluene with Boron Fluoride-t-Butyl Chloride Complex. t-Butyl chloride (0.03 mole) was dissolved in 0.06 mole of toluene, then 0.03 mole of boron fluoride was absorbed completely in this solution a t -95'. When the solution was allowed to warm up to room temperature, gas evolved vigorously. The gas was sampled and analyzed with the mass spectrometer.
Results The melting points of boron fluoride-isopropyl chloride and -t-butyl chloride 1: 1 addition complexes were -136 and -87'. The melting point of the boron fluoride-methyl fluoride complex was - l l O o as observed by Olith, et czL2 These complexes were colorless (by Olith, et al., the colors of boron fluoridepropyl fluoride and -&butyl fluoride complexes are both yellow) and dissociated by reversible reaction into their respective components when the temperature was raised, while colorless boron fluoride-ethyl fluoride complex produced a brown-colored polymer when left for a long time at room temperaturee6 The values of a single-stage separation factor at -112' for several mole ratios of isopropyl chloride to boron fluoride in the liquid phase are shown in Table I. When the mole ratios are equal to or larger than 1, the value of the factor is 1.021, a constant independent of the mole ratio. However, when the mole ratio is smaller than 1, the values become smaller and the plot
Table I : Single-Stage Separation Factor of Boron Fluoride-Isopropyl Chloride System Mole ratio of isopropyl chloride t o boron fluoride in the liquid phase
Temp., "C.
Vapor press.. mm.
Single-stage separation factor,
0.25 0.50 0.75 1 .oo 1.oo 2.00 2.00 3.00 2.00 3.00 3.00
-112 -112 -112 -112 -112 -112 - 112 - 112 - 95 - 95 - 95
260 225 190 160 160 90 90 60 310 250 250
1.013 i 0 , 0 0 1 1.016 f 0.001 1.019 f 0,001 1.021 f 0.001 1.021 f 0,001 1.022 f 0,001 1.021 f 0.001 1.021 f 0.001 1.016 f 0,001 1.017 f 0.001 1.016 f 0.001
LI
(6) R. Nakane, S. Isomura, and 0. Kurihara, presented a t the 3rd Stable Isotope Congress, Leipzig, Germany, Oct. 30, 1963. (7) W. F.Edgell and L. Parts, J . Am. Chem. SOC.,7 7 , 4899 (1955). (8) R. Nakane and 0. Kurihara, Sci. Papers Inst. Phys. Chem. Res. (Tokyo), 5 6 , 161 (1962).
Volume 68, Number 10 October, 1964
R. KAKANE,0. KURIHARA, AND A. KATSUBORI
2878
1.00
0.8
0
0
0
0
0
1.030
d 1.026
s
I
0
.-
2 1.020 8 8
(4)
1.015 0
[B' I F 3 'i-C3H,C1(1i q ) ]
#
=
a' [BF, .i-C3H,C1(1iq) ]
(5) in which a is nearly equal to a', for in a single-stage separation the change of isotopic abundance ratio is very small as given in Table I. Consequently
F z 1.010
1.005
0
0.5
1.0
[E]
1.5
2.0
2.5
3.0
in liquid phase.
Figure 1. Single-stage separation factor vs. mole ratio of alkyl halide to boron fluoride in the liquid phase: 0 , CHaF at -95"; 0, CH3F a t -112"; A, i-C3H&1 a t -95"; A, i-C3H&1 at - 112"; 0 , tCdHyCl a t -112'.
of the mole ratio us. separation factor gives a straight line (Fig. 1). This result can be explained as follows. By mass spectrometry, the gas was observed to be composed mostly of boron fluoride at -95 or - 112'. Therefore, when in the liquid phase no free liquid boron fluoride exists, the isotopic exchange occurs only between gaseous boron fluoride and the liquid complex and the equilibrium constant of this isotopic exchange is given by
However, if free liquid boron fluoride coexists with the complex, the isotopic exchange between gaseous boron fluoride and liquid boron fluoride occurs simultaneously and the equilibriuin constant of this isotopic exchange is given by
The equilibrium constant K 2 is known to be 1.011 a t -l12°.4J I n this system two isotopic exchanges must be involved and, hence, the obtained value of the singlestage separation factor a is given as a function of equilibrium constants K 1and K 2of both reactions as The Journal of Phgsical Chemistry
=
Kz
+ (Kl - K 2 ) x
Since K I and K2 are constants, the plot of a us. [BF,. iC3H7Cl(lls)I/( I B F ~ ( 1I ~ ~ ) [BF3.i-C3H7Cl(llq)I) is a straight line and, when no free liquid boron fluoride exists, a is equal to K1. The obtained result shows the following facts. One mole of boron fluoride fornis with one mole of isopropyl chloride the 'liquid 1: 1 addition complex. When the mole ratio of isopropyl chloride to boron fluoride in the liquid phase is 1,the liquid phase is only of the complex. When it is larger than 1, the complex is dissolved into excess liquid isopropyl chloride. In these two cases, no free liquid boron fluoride exists and, thus, a obtained experimentally is equal to K1, a constant; but when the rnole ratio is smaller than 1, free liquid boron fluoride coexists with the complex in the liquid phase. I n this case, the mole ratio of isopropyl chloride to boron fluoride in the liquid phase becomes equal to the mole fraction of complex in the liquid phase, [BF3.i-C3H&l(lla)]/ ([BF3.i-C3H7C1(1,,,] [BF3(llq)]), for one mole of isopropyl chloride forms with one mole of the boron fluoride complex. Therefore, a plot of a us. mole ratio becomes a straight line. Thus, it is evident that the equilibrium constant of isotopic exchange between gaseous boron fluoride and the liquid boron fluoride-isopropyl chloride complex is 1.021 a t - 112O. By a similar method, the equilibrium constant at'-95O was found to be 1.016, which is given in Table I.
+
+
B O R O N ISOTOPE EXCHAKGE l?ETWEEK
BE',
AND ITS
ALKYLHALIDE CORIPLEXES
The values of single-stage separation factor a t -95' for several mole ratios of methyl fluoride to boron fluoride are shown in Table 11. When the mole ratio in the
Table 11: Single-Stage Separation Factor of Boron Fluoride-Xlethyl Fluoride System Mole ratio of illethyl fluoride to boron fluoride Liquid phase Gas phase
0.25
0.002
0.33
..,
0.38 0.43 0.48 0.53 0.60 0.32 0.40 0.50 0.60 0.80 1.00 1.oo 1.25 1.75
0.006
a
...
0.010 0.016 , . .
0.002 0.009 0.014 0.03 0.13 0.35 0.35 0.80 2.99
Temp., o c .
-112 -112 -112 - 112 - 112 -112 - 112 - 95 - 95 - 95 - 95 - 95 - 95 - 95 - 95 - 95
Vapor press., mni.
Single-stage separation factor, a
205 180 160 140 125 105
1 1 I 1 1 1
014 016 017 018 018 019
i0 f0 f0 f0 i0 f0
001 001 001 001 001 001
680 ,530 420 330 200 140 140 110 110
1 1 1 1 1 1 1
011 012 013 014 017 020 019 1 020 1 020
i0 i0 f0 f0 f0 i0 f0 i0 f0
001 001 001 001 001 001 001 001 001
2879
the separation factors become measurable. The results are shown in Tables I1 and 111. The plots of QI us. mole ratio are also straight lines. I n the boron fluoride-f-butyl chloride system, the boron fluoride molecule alone is found by mass spectrometry to exist in the gas phase. Therefore, in these cases, two isotopic ex-
Table 111: Single-Stage Separation Factor of Boron Fluoride-&Butyl Chloride System Mole ratio of t-butyl chloride to boron fluoride in liquid phase
0.15 0.25 0.30 0.38 0.45 0.48 0.50 a
Temp., O C .
-112 - 112 - 112 - 112 -112 - 112 - 112
Single-stage separation factor, a
Vapor press., mm.
285 275 270 260 250 245
1.015 i 0,001 1.016 f 0.001 1.018 f 0.001 1.019 i 0.001 1.022 f 0.001 1.021 f 0.001
...
...
Solid.
Solid
liquid phase is equal to or larger than 1, the values become constant, but when it is smaller than l , the plot of a us. mole ratio becoines a straight line as in the case of the boron fluoride-isopropyl chloride system. It was observed by mass spectrometry that even in the gas phase, methyl fluoride molecules coexisted with boron fluoride rnolecules. If both molecules can form a gaseous complex, the isotopic exchange between gaseous and liquid complexes should occur together with that between the gaseous boron fluoride and liquid complex and, thus, the value of the separation factor should depend also on the molc ratio in tho gas phase, but such a positive relation between the value of a and the mole ratio in the gas phase could not be obtained. Therefore, in the boron fluoride-methyl fluoride system, too, the boron fluoride-methyl fluoride complex exists only in the liquid phase, and for the complex, the isotopic exchange between the gaseous boron fluoride and liquid complex alone occurs at low teinperat>ures with the value of equilibrium constant 1.020 at -95'. Boron fluoride-inethyl fluoride and -&butyl chloride complcxes are both solid at - 112', but when the mole ratio of inethyl fluoride or t-butyl chloride to boron fluoride in the liquid phase is smaller than 1, solid complexes are dissolved into liquid boron fluoride and
changes between gaseous boron fluoride and liquid complexes and between gaseous and liquid boron fluoride occur simultaneously. The values of the separation factor, which are obtained a t the mole ratio 1 by extrapolation of the straight lines, should become the equilibrium constants of these complex systems. Thus we obtain the equilibriuin constant of the boron fluoride-methyl fluoride system as 1.026 and that of the boron fluoride--t-butyl chloride system as 1.033 both at -112O. Alkylation of Toluene b y the Boron Fluoride-t-Butyl Chloride Complex. When boron fluoride gas was absorbed in a toluene solution of t-butyl chloride at low temperatures, the soIution becomes almost colorless. When this solution was allowed to become warm, there was a vigorous gas evolution. Mass spectrometrical analysis of the evolved gas showed the production of hydrogen chloride and the disappearance of t-butyl chloride in this reaction, resembliiig the alkylation of alky benzene with alkyl fluorides except for the color change in the inter~nediate.~ The preceding paragraph shows for the first time that t-butyl chloride can readily alkylate toluene in the presence of boron fluoride, and the common assertion that boron fluoride cannot catalyze the alkylation of aromatic compounds with alkyl chloride is now denied, sxcepting the cases of methyl and propyl chlorides. Volume 68, Sumher IO
October. l N 4
R. XAKANE, 0. KURIHARA, AND A. NATSUBORI
2880
Discussion The values of equilibrium constants of isotopic exchange between boron fluoride gas and many boron fluoride complexes a t various temperatures are surnmarized in Table IV. The equilibrium constant is the
Table IV : Equilibrium Constant of Boron Isotope Exchange between Boron Fluoride Gas and Liquid Boron Fluoride Complexes Liquid complex
-Equilibrium 25' -95'
1.028 (1 .071)a,a ... 1.031 . . .a 1.032 . . .a 1.027 . . .0 1.033 ... 1.023 . . .a 1.036 ... 1.026
1.020 1.016 1.015 1.013 1.012
1.006 ( 1 .006)b
constant,, K-96.5'
-112'
Ref.
(1 .082)a,b c . . .a d . . .a e . . .a e ... e . . .D f ... f . . .0 d 1 .033a 9 1.022 ...a h 1,026" g 1.021 g 1.016 1.021 h 1.020 i 1.018 j 1.017 i 1.011 1.017 h 1.011 j 1.011 i 1.011 h, j
a Solid. Extrapolated from experimental equation. See S. V. Ribnikar, "Proceedings of the International ref. 9. Symposium on Isotope Separation," North Holland Publishing Co., Amsterdam, 1958, p. 204. e R. hf. Healy and A. A. Palko, J . Chem. Phys., 28, 211 (1958). f S. V. Ribnikar and C . A. Bootsma, Bull. Inst. Nucl. Sci., "Boris Kidrich" (Belgrade), 9, 91 (1959). Present report. h See ref. 8. See ref. 5. 1 See ref. 4.
largest for the system of strong complex that is stable even at room temperature; it is the smallest for the system of liquid boron fluoride or boron fluoride in nonpolar monoolefin solution in which an unstable r-complex is formed. The system of weak complexes that are stable only a t low teniperatures has a n intermediate value. This suggests that the value of the equilibrium constant relates to the stability of the complex. Concerning the values of the equilibrium constants of systems of strong coinplexes, the two following points are noticeable. First, they all are approximately 1.03 a t rooni temperature as shown in Table I V regardless of the sorts of donor molecules, and they may roughly be considered equal to each other when they are compared T h e Journal of Physical Chemistry
with the values of the constant of systems of complexes of other types. Second, Palko, et al., worked on the isotopic exchange between boron fluoride gas on one side and boron fluoride-methyl ether,g-ethyl ether, and -tetrahydrofuran'" complexes on the other, and found that only the isotopic shifts of two B-F antisymmetric stretching frequencies, which became lower by transformation of boron fluoride from planar to tetrahedral, had a very large effect upon the partition function ratio between isotopic complex molecules, and the contribution of the other vibrations such as B-0 stretching to the ratio was much smaller. I t is well known" that in complex formation of boron fluoride with methyl ether, the boron and oxygen atoms are tetrahedrally bonded with a lone pair of electrons occupying the fourth orbital of the oxygen atom and the boron atom is situated a t the center of the tetrahedral XY3Z form, the symmetry of which is CaY. I t seems likely that in every molecule of all strong complexes boron valence bond angles are approximately tetrahedral, the B-F distancc is presumably nearly equal to that of boron fluoride-methyl ether complex, and hence the B-F antisymnietric stretching frequencies of all these complexes are nearly equal to one another. For example, Palko, et al., observed'" that the B-F antisymnietric stretching frequencies of B"F3 (CH,)20,R11F3.(C2H5)20, and B1'F3.(CH2)40are nearly equal. Waddington, et a1.,I2 reported also that those of BF3~POC13, BF3.(C2H&0, and BF3Cl- (in PH4BF3C1, NOBF3C1, and (CH3),KBF3C1) except BY3.CjH5iY, are nearly equal. On the other hand, on other tetrahedral XY3Z molecules (C3") such as methyl halidesI3 or silyl halides,14 it is known that X-Y antisymmetric stretching frequency does not change when the Z atom is replaced by another. That the isotopic shift of antisymmetric stretching frequency is much larger than any of the isotopic shifts of other vibration frequencies is also found on tetrahedral XY4 molecules ( T d ) , in which the Y atom is not hydrogen. When the center atom X in CF415or BF4-,I6 (9) A. A. Palko, G. 552 (1962).
M . Begun,
and L. Landau, J . Chem. P h y s . , 37,
(10) G. M. Begun and A. A. Palko, ibid.,38,2112 (1963). (11) 5. H. Bauer, G. R. Finlay, and -4.W. Laubengayer. J . Am. Chem. SOC.,65, 889 (1943); 67,339 (1945). (12) T. C. Waddington and F. Klanberg, J . Chem. SOC.,2239 (1900). (13) G. Heraberg, "Molecular Spectra and Molecular Structure. 11. Infrared and Raman Spectra of Polyatomic Molecules," D. Van Nostrand, New York. N. Y., 1945, p. 135. (14) C. Newman, J. K. O'Loane, 5. R. Polo, and M .K. Wilson, J . Chem. P h y s . , 25, 855 (1956). (15) J. Goubeau, W. Bues, and F. W. Kampmann, 2. anorg. allgem. Chem., 283, 123 (1956). (16) J. Goubeau and W.Bues, ibid., 268, 221 (1952); N. S . Greenwood, J . Chem. Soc., 3811 (1959).
BORON ISOTOPE EXCHASGE BETREEK BF3 AND ITSALKYLHALIDE COMPLEXES
for example, is replaced by its isotope, X-Y antisynimetric stretching frequency alone shows a large isotopic shift. Therefore, in all strong boron fluoride complexes it can be assumed that all values of the B-F antisymmetric stretching frequency are nearly equal and that the effect of isotopic shift upon the partition function ratio between isotopic complex molecules is far larger than that of other vibration frequencies. If so, the equilibrium constant will be nearly the same. The fact that all the observed values are approximately 1.03 at room temperature proves the assumption mentioned above being reasonable. Thus, for all the systems of a strong complex, the equilibrium constant will be approximately 1.08 a t - - 1 1 2 O , which agrees with the value 1.082 a t - 112' calculated from the empirical formula given by Palko, et al., for the boron fluoridemethyl ether system. I n an uncoordinated boron fluoride molecule, too, the isotope shift of the doubly degenerate B-F antisymmetric stretching frequency is the largestI7 and their effect upon the partition function ratio between isotopic boron fluoride molecules is much larger than that of the other frequencies. Therefore, it can be assumed, although roughly, that in boron fluoride complexes of all the types, B-Z stretch contributes only slightly to the partition function ratio and the isotopic shifts of B-F antisymmetric stretching frequencies alone have a very large effect on it. If so, the partition function ratio between isotopic complex molecules, Q [B"F3. ZA]/ Q [B10F3.2-41,becomes smaller and therefore the equilibrium constant
&(B"I'd Q(B"F3)
/
Q[B"F,.ZAl Q [B1"F3.ZA]
becomes larger as the B-F antisymmetric stretching frequency of the complex decreases. Thus, the observed result that the value of equilibrium constant of the system of weak complex is between the values of the constant of liquid boron fluoride and strong boron fluoride complex shows that the B-E' antisymmetric stretching frequency of weak boron fluoride complex lies between the frequencies of uncoordinated boron fluoride and strong boron fluoride complex. As the density OS electrons localized on the vacant orbital of boron atoms increases, more boron fluoride molecules transform from planar into tetrahedral and the B-F antisymmetric stretching frequency becomes lower, although the density of localized electrons, which determines the strength of the B-Z bond, will be only indirectly related to the B-F stretch. Since the equilibrium constant depends mostly on the B-F antisymmetric stretching frequency as mentioned above, it follows that the density of localized electrons on the
2881
vacant orbital of boron atoms and hence the polarity of the complex becomes higher as the equilibrium constant becomes larger, causing the complex to become more stable. Now, since the equilibrium constants obtained for boron fluoride-alkyl halide systems have intermediate values, all these boron fluoride-alkyl halide complexes will be weak polar complexes. Therefore, the boron fluoride-t-butyl chloride complex seems to be a polar 6-
6T
complex, BF3+C1C(CH3)3, but not a carbonium ion complex, kF,Cl.&(CH,), although t-butyl chloride is easily ionized. The polarity of these complexes will 6-
6+
6-
6+
6+
6-
decrease in the order: BFp-C1C(CH3)3> BF3+FCH3
>
6-
6+
BF,+ClCH(CH,),, BF3+ClCH3. If so, the boron fluoride-t-butyl chloride complex too should have catalytic activity similar to the boron fluoridemethyl fluoride complcx. This presumption was proved to be correct, for t-butyl chloride can alkylate toluene readily in the presence of boron fluoride, although such a color change as observed by 01&h,et aL,*on alkylation with the boron fluoride-alkyl fluoride complex was not found. However, despite the alkylation of aromatic constituent with tertiary alkyl halide, it seems very likely that the alkylation does not proceed through the formation and reaction of the carbonium ion, for this complex is considered to be only a weak polar complex as mentioned above. Thus, the reaction is presumed to be that a stable intermediate colorless a-complex is produced through an activated complex in a transition state and then it decomposes to alkyl benzene, hydrogen chloride, and boron fluoride by the process
6 + y -6 -CI-BF,
....H
6-
(CW3
C -.-CI-BF3
I
(CH3)3
(activated complex)
'iH3
(17) J. Vanderryn, J . Chem. Phys., 30, 331 (1959).
V o l u m e 68, 'V'umber 10
October, 1064
2882
AMOSJ. LEFFLERAND
On the other hand, the boron fluoride-isopropyl chloride complex, the equilibrium constant of which system is smaller than that of the boron fluoride-methyl fluoride system, cannot alkylate toluene despite being also polarized. It is interesting to compare generally the equilibrium constants of boron fluoride complex systems with the catalytic activity of the boron fluoride complex. Boron fluoride-ether and -alcohol complexes, which are cornpletely ionized and have large equilibrium constants, are very active catalysts. On the other hand, the boron fluoride-t-butyl chloride complex, which has a small value of equilibrium constant, is a less reactive catalyst but can alkyIate toluene readily. However, the boron fluoride-alkyl halide complex, the equilibrium
SORMAX
11. \+'IEDERHORN
constant of which is smaller than 1.021 a t -112O, cannot have even the catalytic activity for the alkylation. As shown in the previous w o r k ~ , boron ~J fluoride polar monoolefin a-complex systems have such values of equilibrium constant as approximately 1.02 at - 112 O, yet none of the complexes can react with free olefin monomer molecules, and only when there are strong complexes, e.g., B F 3 . H z 0 , whose systems have large values of equilibrium constant, does the polymerization of monoolefin proceed. Therefore, in order that boron fluoride complexes have any catalytic activity, the electrons from donor molecules must remain localized on vacant orbitals of boron atoms so that the value of the equilibrium constant becomes larger than approximately 1.03 a t -112" and the catalytic activity beconies higher as the equilibrium constant increases.
Thermodynamics of the Liquid Potassium-Oxygen and Sodium-Oxygen Systems
by Amos J. Leffler and Norman M. Wiederhorn Arthur D. Little, Inc., Cambridoe, Massachusetts
(Received February 7 , 196'4)
The oxygen pressure-melt composition for the potassium-oxygen system was observed at three different temperatures. From these data and the tabulated thermodynamic values for KzO, the heats and free energies of formation of the melt compositions were calculated from KOo5to KOz.o,although experimentally it was not possible to reach conipositions richer in oxygen than KO1 75, A study was made of the oxygen pressure-melt composition for the liquid sodium-oxygen system between 780 and 980'. These results together with the therinodynaniic data for NazOzwere used to calculate the heats and free energies of formation of melt compositions between NaOl and NaOo 6.
Introduction The properties of the oxides of potassium have been partially investigated and thermodynamic values have been tabu1ated.l The most extensive experimental data are available for KO2, including heats of formationJ2- specific heats from 20 to 373°K.,5,6 and heats and free energies of decompositioni8 to Kz02a t 300-370° in the solid phase. An experimental heat of formation for K 2 0 3 has been determined. The only available The J o u r n a l of Physical Chemistry
information on KzOzis the heat and free energy of decompositioni in the 300-370' range. During a crystallographic study of KOzl" a phase transition was noted (1) J. F. Coughlin, U. S. Bureau of Mines, Bulletin 542, U. S. Covt. Printing Office, Washington, D. C., 1953. (2) P.W.Gilles and J. L. Margrave, J . P h y s . Chem., 6 0 , 1333 (1956). (3) L. J. Kazarnovskaia and I. A. Kazarnovskii, Z h . Fiz. K h i m . , 2 5 , 293 (1951). (4) R. de Forcrand, Compt. rend., 158, 991 (1914).
