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of cobalt hydrocarbonyl with cyclohexene5 apparently occurred only when the temperature of the mixture had risen to a point (15") where cobalt hydrocarbonyl in the concentrations used is known to decompose rapidly. l1 We believe these results are sufficient to warrant a re-examination of the fundamental postulate as to the nature of cobalt catalysis of the hydroformylation reaction. It appears likely that any direct role in hydroformylation catalysis that cobalt carbonyl or cobalt hydrocarbonyl may play must be carried out in conjunction with a solid cobalt surface.
tems with diborane and the following ethers: diethyl ether, methyl ethyl ether, dimethyl ether, ethylene oxide, tetrahydropyran, tetrahydrofuran, perfluoroether [ (CzF&O] and Cyclo-CIFsO.
the diborane pressure, suggesting the formation of a tetrahydrofuran-borine complex in solution. No indication of a diethyl ether-borine complex was found. The existence of the complex C4HsO :BH3 was established definitely by Rice, Livasy and Schaeffer4 from the solid-liquid equilibrium in the
Results and Discussion Diethyl ether (Fig. 1) forms two complexes with diborane, the expected borine complex (CzH&O : BH3 and a second complex having the composition (C2H5)20.3BHa. Methyl ethyl ether (Fig. 2) does not form a 1:1borine complex, but does form a congruent melting compound CH30C2H6.2BH3.
Experimental Materials.-Diethyl ether, tetrahydropyran and tetrahydrofuran were dried over lithium aluminum hydride and fractionally distilled before use. Methyl ethyl ether was prepared by the reaction of ethyl iodide and sodium methylate in absolute methanol and purified by fractional distillation from lithium aluminum hydride at 5". The above ethers, as well as dimethyl ether, perfluoroether, cycloCdFsO,e ethylene oxide and diborane (su plied by .the Olin Mathieson Chemical Corporation) was Lrther purified if necessary by fractional distillation a t low temperature (11) H.W.Sternberg, I. Wender, R.A.Friede1 and M. Orchin, J . Am. ' was attained, as until a purity of better than 98 mole % Chem. Soc., 7 6 , 2717 (1953). determined by freezing point measurements. Apparatus.-A freezin point cell of the type described by Davidson, Sisler and [toenner? was attached to a vacuum COMPLEXES OF ETHERS gas-handling line which included manometers, calibrated flasks and traps for low temperature distillation. The WITH DIBORANEl composition of the liquid mixtures used was calculated BY HENRY E. WIRTH, FRANKLIN E. MASSOTHAND DAVID from the pressure, volume and temperature of each of the X. GILBERT gaseous components. From the volume of the freezing point cell and the observed pressure, the total moles of diDepartment of ChemiStTtt, Syracuse University, Syracuse, N . Y . borane added were corrected for the moles present in the Received January l i s 1968 gas phase at the freezing point of the mixture, assuming all to be diborane. I n 1938, Schlesinger and Burg2 showed that di- gasApresent copper-constantan thermocouple was used to deterborane and dimethyl ether formed the complex mine temperature. Cooling curves were recorded by a (CH3)20:BH3at -80". Elliot, Roth, Roedel and Brown Electronik recorder (full scale = 1 mv., variable Boldebuck3 found that the solubility of diborane range). Temperatures were reliable to =k0.5", and mole in tetrahydrofuran depended ,on the square root of fraction to f2%.
1
IbO I -
90
100 0
I
I
10
Po
II 30
I
I
I
1
I
I
J
40
60
60
TO
00
94
100
Mole 'lo Diborane.
Fig, 1.-The
diethyl ether-diborane system.
diborane-tetrahydrofuran system. They suggest that the relative stability of the ether-borines is in the order C4HsO:BHz > (CHa)zO:BH3 >> (CzHs)zO:BH,
This order of stability was confirmed by Raman ~pectra.~ In this work phase diagrams are reported for sys(1) This research was supported in part by the Department of the Navy, Bureau of Aeronautics, through subcontract with the Olin Mathieson Chemical Corporation. (2) H. I. Schlesinger and A. B. Burg, J . Am. Chem. Soc., 60, 290 (1938). (3) J. R. Elliot, W. L. Roth, G . F. Roedel and E. M. Boldebuck, ibid., 74, 5211 (1952). (4) B. Rice, J. Livasy and G . W. Schaeffer, ibid., 77, 2750 (1955). 15) B. Rice and H. 8. Uchida, THISJOURNAL, 59, 650 (1955).
o
I
I
I
I
1
I
I
I
I
l
IO
20
so
40
so
eo
70
80
90
100
Mole "1. D i b o r o n e .
Fig. 2.-The
methyl ethyl ether-diborane system.
Ethylene oxide (Fig. 3) forms no congruent melting compounds, but the freezing point diagram indicates the possibility of one and perhaps two incongruent melting compounds. The initial freezing points and the "flats" at 131 and 121°K. were definite and reproducible, but the secondary breaks between 70 and 95 mole yo diborane were erratic. The formation of solid solutions in this range could explain these observations. While ethylene oxide is known t o react vigorously with diborane at 190°K.,8 there was no evidence of reaction at tem(6) The fluorinated ethers were supplied to us through the courtesy
of the Minnesota Mining and Manufacturing Co. (7) A. W. Davidson, H. H. Sisler and R. Stoenner, J . A m . Chem. Soc., 6 6 , 779 (1944). (8) F. G . A. Stone and H. J. Emeleus, J . Chem. SOC.,2755 (1950).
NOTES
July, 1958
87 1
CIHsO:BF3, (CH&0:BF3, (CzH&0:BF3, (iso-CsHT)z:BF3
100
O
I
I
10
10
1 ;
I
30
40
50
I
I
I
I
*O
IO
W
80
100
Mole % Diborane.
Fig. 3.-The
ethylene oxide-diborane system.
peratures slightly above the melting points of the mixtures. Dimethyl ether, tetrahydrofuran and tetrahydropyran (Figs. 4A, 4B and 4C, respectively) all form 1:1 borine complexes. The whole range of compositions could not be investigated, as the diborane pressure exceeded one atmosphere when the mole fractions of diborane were above 0.3-0.4.
This order cannot be explained on the basis of the inductive effect of the alkyl groups, but is accounted for by the effect of steric strains. The same order of relative stability was found, as expected, for the first three corresponding borine complexes by Rice and Uchida.6 It would be expected however that a borine complex of methyl ethyl ether would be intermediate in stability between those of methyl and ethyl ether, and that ethylene oxide would form a strong complex. That they do not show up on the phase diagrams may indicate that a t low temperatures in these solvents the concentration of borine from the dissociation of diborane is too low to permit the formation of 1:1 complexes. The inductive effect of the highly fluorinated substituents in perfluoroether and cyclo-C4F80reduces the basic strength of these ethers so they cannot form complexes even with boron trifluoride. The unexpected compounds found with methyl ethyl ether and diethyl ether could be regarded as molecular addition compounds of diborane t o which the formulas CH30C2H5.B2H,and (CzHb)20: BH3. BzHe could be applied: By analogy to the corresponding boron trifluoride complexes, lo the formulations CH30C2H6.2BH3and (C2H5)20.3BH3are preferred. (IO) H. E. Wirth, M. J. Jackson and H. W. Gritlitha, THISJOURNAL, 62, 871 (1958).
COMPLEXES OF ETHERS WITH BORON TRIFLUORIDE 2
OL
I 2 00
I
I
I
10
20
30
BY HENRYE. WIRTH, MIRIAMJ. JACKSON A N D HOWARD W. GRIFFITHS 40
0
30
20
IO
40
Department of Chemistry, Syracuse University, Syracuse, AT. Y. Received January 1 1 1968 ~
I00
140
- 1I4 0 0
10
20
30
40
0
20
Mole % D i b o r a n e .
40
eo
no.
__
ieo
Fig. 4.-A, dimethyl ether-diborane; B, tetrahydrofurandiborane; C ,tetrahydropyran-diborane; D, perfluoroetherdiborane.
Perfluoroether and diborane (Fig. 4D) formed a partially miscible liquid system, with no indication of compound formation. Cyclo-C4F80 formed a similar system. Two liquid phases also were observed when perfluoroether and boron trifluoride were mixed a t 155°K. Brown and Adamss have shown that the relative stability of the etherates of boron trifluoride decreases in the order: (9) H. C. Brown and M. Adams, J . A m . Chem. SOC.,64, 2557 (1942).
Although 1:1 ether-boron trifluoride complexes are well known there are no reports in the literature of compounds containing more than one mole of boron trifluoride per mole of ether except for a single reference’ to a second complex of diethyl ether containing 60-90 mole % BF3. I n their monograph Booth and Martin2 refer to the compounds 2(CH3)2O.BF3 and 2(CzH&0 .BF3, but the reference cited3 discusses compounds of the alcohols. Since recent work4 in this Laboratory has demonstrated the existence of the borine complexes (CzHs)20.3BH3and CH30C2H6.2BH3, a survey of several ether-boron trifluoride systems was made to see whether similar compounds are formed with boron trifluoride. Experimental The apparatus described in the previous article4was used. Methyl n-propyl ether and methyl isopropyl ether were prepared by the reaction of sodium methylate with 1bromopropane and 2-bromopropane, respectively, in methyl alcohol solution. The mixtures were refluxed for 3 hours, and the material distilling below 60” collected. Thme products were washed with water, dried successively over (1) A, F. 0. Gerrnann and M. Cleaveland, Science, 63, 582 (1921). (2) H. S. Booth and D. R. Martin, “Boron Trifluoride and Its
Derivatives,” John Wiley and Sons, Inc., New York, N. Y., 1949. (3) H. Meerwein and W. Pannwitz, J . prakt. Chem., 141, 123 (1934). (4) H. E. Wirth, F. E. Massoth and D. X . Gilbert, TIUS JOURNAL^ 62, 870 (1958).
