ELECTRICAL CONDUCTIVITIES OF BF3 IN HALOGEN FLUORIDES
224 1
The Electrical Conductivities of Boron Trifluoride in Pure and Mixed Halogen Fluorides’
by Madeline S. Toy and William A. Cannon Douglas Missile and Space Systems Division, Astropower Laboratory, Newport Beach, California (Received January 6 , 1966)
The electrical conductivities of solutions of boron trifluoride in chlorine and bromine trifluorides have been studied as a function of temperature and concentration. Dilute solutions of boron trifluoride in bromine trifluoride have conductive properties similar to strong electrolytes in water, with an approximate Onsager slope of 80.2 and equivalent conductance at infinite dilution of 131 cm2ohm-’ equiv-’ at 25’, while the equivalent conductances of dilute solutions of boron trifluoride in chlorine trifluoride decrease with dilution. Their differences and mechanisms are discussed. The preparation and identification of difluorobrominium tetrafluoroborate is described.
Introduction The physical and chemical properties of halogen fluorides have been studied extensively. Particular interest has been shown in the self-ionization of many of the halogen fluorides. Haszeldine3 has observed that chlorine trifluoride dissolves in both bromine trifluoride and iodine pentafluoride and is not evolved upcn heating the resulting solution to 60-70”, although chlorine trifluoride boils at 12”. Whitney and coworkers4have treated alkali metal fluorides in chlorine trfluoride and bromine pentafluoride a t 100’ to form the respective alkali metal tetrafluorochlorates and hexafluorobromates; the compounds KCIF4, RbCIF4, CsC1F4, KBrF6, RbBrF6, and CsBrFs have been prepared. Rogers and Katz6have studied the exchange of radioactive F’* between chlorine and bromine trifluorides and suggested that chlorine trifluoride acts as a base in bromine trifluoride solution, the exchange reaction proceeding through the equilibrium C1F3
+ BrF3 ZC1F2++ BrF4-
(1)
The electrical conductivity of mixed halogen fluorides has been reported by Quarterman, Hyman, and Katz6 for solutions of bromine trifluoride and bromine pentafluoride. Boron trifluoride itself is a nonconductor of electricity. The solutions of boron trifluoride and halogen fluorides are examples of electrolytic conduction
caused by “potential electrolytes” in the sense of the definition given by Kortum and Bockris.’ The nonpolar BF3 molecules react with the self-ionizing BrF3 molecules to form electrolytes as described by eq 2 and 3. BFdg)
+ BrF3(1) 1-BrF2+ + BFb-
(2)
+
2BrF3(1) BrF2+ BrF4(3) In the case of BF3 in a mixed solvent system (BrF3ClFa), the set of eq 1 to 5 is postulated. BF3(g)
+ ClF3(1) ZClFz+ + BFI-
(4)
+
2CIF3(1) C1F2+ ClF4(5) Selig and Sham? have prepared difluorochlorinium tetrafluoroborate according to eq 4, and we have pre(1) Presented in part at the Western Regional Meeting of the American Chemical Society, Los Angeles, Calif., Nov 20, 1965. (2) A. G. Sharpe, Quart. Rev. (London), 4,115 (1950); V. Gutmann, Angew. Chem., 62, 312 (1950); H. C. Clark, Chem. Rev., 58, 869 (1958). (3) R. N. Haszeldine, J . Chem. SOC.,3037 (1950). (4) E.D.Whitney, R. 0. MacLaren, C. E. Fogle, and T. J. Hurley, J . Am. Chem. SOC.,86, 2583 (1964). (5) M. T.Rogers and J. J. Katz, ibid., 74, 1375 (1952). (6) L. A. Quarterman, H. H. Hyman, and J. J. Katz, J . Phys. Chem., 61, 912 (1957). (7) G. Kortum and 0. Bockris, “Textbook of Electrochemistry,’’ Vol. I, Elsevier Publishing Co., New York, N. Y.,1951,pp 99,100. (8) H. Selig and J. Shamir, Inorg. Chem., 3, 294 (1964).
Volume 70,Number 7 July 1066
2242
MADELINE S. TOYAND WILLIAMA. CANNON
pared difluorobrominium tetrafluoroborate according to eq 2 (see Experimental Section). For the isolation of a pure and stable product, a lengthy evacuation is required to eliminate trace amounts of BrF3 (bp 125127") or compounds as zBrF3.yBF3trapped in the pure BrF3.BF3solid. Experimental Section
Materials. Boron trifluoride, bromine trifluoride, and chlorine trifluoride were obtained from the Rfatheson Co. Chlorine trifluoride and boron trifluoride were purified by passing the vapor through a sodium fluoride scrubber to remove possible hydrogen fluoride impurity. Bromine trifluoride was purified by removing the first fraction under reduced pressure a t room temperature. The melting point and specific conductivity of the bromine trifluoride agreed with the literature value^.*^^^^ The purity of boron trifluoride and chlorine trifluoride was checked by vapor spectroscopic spectra and specific conductivities of the liquids. These samples were indistinguishable from the samples which had been exhaustively fractionated. However, the specific conductivity of chlorine trifluoride employed, 8.8 X ohm-' cm-', was slightly higher ohm-' cm-1.2 than the best literature value, 3.9 X Apparatus. The conductivity cells were made of borosilicate glass with smooth platinum electrodes. Cell constants ranged from 0.03 to 0.5 cm-'. An internal thermocouple well located near the electrodes was used for temperature measurement. Cell-resistance measurements were made with a General Radio Type 1650-A impedance bridge. It was equipped with an internal, 1000-cycle signal source and tuned null detector. Measurements were also made with an external signal source ranging from 200 to 10,000-cycles to permit correction for polarization. For more sensitive balance a Hewlett-Packard 400 L vacuum tube voltmeter was used as an external null detector. A stainless steel vacuum line constructed of 304 stainless steel tubing and 316 stainless steel needle valves was employed for manipulation of boron trifluoride and halogen fluorides. Mercury manometers cannot be used; therefore pressure was measured with a stainless steel Bourdon gauge covering the range from 30-in. vacuum to 30 psi pressure. Careful exclusion of moisture is essential in handling these compounds as the presence of any water vapor will cause noticeable attack on glass and stainless steel and will likewise affect conductivity measurements. The metal parts of the vacuum system were first degreased with trichloroethylene, rinsed with Freon, then vacuum dried. After assembly, the vacuum system was passivated by filling The Journal of Physical Chemistry
o B4'F3
0:z
0.4 0. 6 0.8 Mol Fraction (N)
1.0
c1F3
Figure 1. Specific conductivities of solutions of BrF3 and CIFs a t 0".
with gaseous fluorine to 1 atm for 18 hr. Finally, the system was evacuated. Under these conditions, it is found that bromine trifluoride does not appreciably attack the stainless steel or glass during the time required to carry out the experiments, viz. 1 to 2 hr. The specific conductivity of chlorine trifluoride is so low that its conductivity is altered by trace impurities from attack on glass. This could readily account for the difference in the observed specific conductivity reported here and the lowest value reported in the literature.2 Preparation of DiJluorobrominium TetraJEuoroborate. The adduct was prepared by bubbling gaseous boron trifluoride through liquid bromine trifluoride a t room temperature. Rapid initial gas absorption was accompanied by exothermic reaction. The solid product did not precipitate, but upon vacuum evaporation to dryness at room temperature, a white residue, mp (in sealed tube) 180" dec, was obtained. The same solid (9) M. S. Toy and W. A. Cannon, "Symposium on Advanced Propellant Chemistry," Advances in Chemistry Series, No. 66, American Chemical Society, Washington, D. C., in press.
ELECTRICAL CONDUCTIVITIES OF BF3 IN HALOGEN FLUORIDES
2243
Temperature
Temperature (OC) 60
40
0
(OC)
-20
-40
Figure 3. Effect of temperature and concentration on specific conductivity of BFs in solutions of BrF3 and ClFs,
1 OW! Temperature
(OK)
Figure 2. Specific conductivities of solutions of BrF8 and C1F3&s a function of temperature.
was obtained by condensing gaseous boron trifluoride in excess on top of frozen bromine trifluoride at liquid nitrogen temperature, warming to room temperature, and vacuum evaporating to dryness. When gaseous boron trifluoride was introduced into a closed evacuated system above liquid bromine trifluoride a t room temperature, a pressure drop was observed as boron trifluoride dissolved. This process was repeated until no additional gas dissolved at 1 atm. Then the yellow liquid was vacuum evaporated a t room temperature to dryness. The yield of the same white product was poor in all these preparations. The solid sample should always be evacuated at room temperature for several hours, preferably overnight, before vacuum sealing the sample for analytical purposes. Anal. Calcd for BrF2BF4: F, 55.69; Br, 39.04; B, 5.33. Found: F, 55.37; Br, 38.71; B, 5.91. The infrared spectrum of thin solid film of the product shows a strong broad band in the region 1020-1100 cm-' indicative of tetrafluoroborate ion.'O A 0.3 M solution of BrF2BF4in BrF3 possesses a specific con-
ductivity of 3.31 X low2ohm-' cm-1 at 25", while the specific conductivity of BrF3 a t 25" is 8.0 X ohm-l cm-'. The formula (BrF2)f (BFJ- is postulated as consistent with conductivity and infrared data.
Results and Discussion Figure 1 shows the conductivities of solutions whose composition lies between pure bromine trifluoride and pure chlorine trifluoride. The conductivity decreases smoothly from the maximum in pure bromine trifluoride to a minimum in pure chlorine trifluoride. The conductivity of pure liquid bromine trifluoride and chlorine trifluoride are almost temperature independent (Figure 2), while the solutions of the two solvents show very slight negative temperature coefficients of conductivity. Figure 3 shows the temperature coefficient of conductivity for BFDin pure BrF3and solutions of BrF3 and ClF3 as similar and both being smaller than aqueous KC1 solutions. Figure 4 shows the increase of conductivity of halogen fluorides with increase of concentration of boron trifluoride. By subtracting the specific conductivity of solvent from the specific conductivity of the solution, the equivalent conductance us. the square root of concentration of difluorobrominium tetrafluoroborate can be estimated (10) R. W. Sprague, A. B. Garrett, and H. H. Sisler, J . Am. Chem. Soc., 82, 1059 (1960); F. See1 and 0. Detmar, 2. Anorg. Allgem.
Chem., 301,8 (1959).
Volume 70,Number 7 J u l y 1966
MADELINE 8. TOYAND WILLIAM A. CANNON
2244
160
A B r F 3 at
I
25OC
BrFj-C1F3 ( 0 . 6:,4,,'b
140
at
Doc
0 BrF3-CIF3 ( 0 . 1 4 NCIF) at O°C
V ClF3 at O°C Aq.KC1 at 25OC
I20 I
'.E 2. I
'ElOo
d
N
uE u 80
B
8
p 60
9a
.h
3 40
0 BrF3 at 25'C 0 BrF3-C1F3Soln.(0. 64NClFJatC
20
A BrF3 -ClF,Soln, (O.74NClFiat' o clF3at OOC
b
I
0.8
I 1.0
a 1
0.0
0.4
Concentration (Mole8 BF3 Per 1000 g Solvent)
Figure 4. Specific conductivity as a function of concentration of BFa in solutions of BrFt and CIF+
and extrapolated to infinite dilution. The approximate equivalent conductance of infinite dilution of BrF2BF4in BrF8 is 131 em2 ohm-' equiv-l with an Onsager slope of 80.2 a t 25". The aqueous potassium chloride curve" is included in Figure 5 for comparative purposes. In dilute concentrations of BF3 in BrF3, the solutions resemble strong electrolytes in water, but BF3 in ClF3 solutions behave very differently. This behavior is in accordance with the presumed high dielectric constant of liquid B r F t and the low dielectric constant of liquid C1F3.12 Thus the phenomenon of decreasing equivalent conductance with decreasing concentration of boron trifluoride in liquid ClFa can be
The Journal of PhysicaE Chem&trg
1.2 1.6 Concentration (Equiv Literm1)
0.8
2.0
2.4
Figure 5. Equivalent conductance of BF3 in solutions of BrFs and CIFs as a function of concentration.
explained by the stable association of ions to form ion pairs as compared with simply solvated ions for high dielectric constant solvents.
Acknowledgment. The authors wish to thank the U. S. Army Advanced Research Project Agency for support of this work under Contract No. DA-31-124ARO-(D)-115. (11) J. F. Chambers, J. M. Stokes, and R. H. Stokes, J. Phya. Chem., 60, 985 (1956); R. A. Robinson and R. H. Stokes, "Electrolyte Solutions." 2nd ed, Butterworth and Go. Ltd., London, 1959, pp 164,466.
(12) M. T.Rogers, H. B. Thompson, and J. L. Speirs, J. Am. Chem. SOC.,76, 4841 (1964).
