Charles B. Colburn
Redstone Arsenal Research Division Rohm & Haas Company Huntsville, Alabama
I
I
Recent Developments in N-F Chemistry
The purpose of this paper is twofo~d: to point out recent developments in the chemistry of compounds containing the N-F bond and to illustrate the absolute necessity of proper instrumentmationin carrying out modern inorganic chemical research. Nitrogen Trifluoride and Fluorine Azide (NFz and FNS
The first fluoride of nitrogen to be prepared was nitrogen trifluoride (NFd, in 1928 by Otto Ruff (1) of Germany after 25 years of experimentation. [Ruff's first reported attempt to prepare a fluoride of nitroren was in 1903 (@.I Nitrogen trifluoride was prepared by the electrolysis of molten ammonium bifluoride. Rnff also reported that difluoramine (HNFz), fluoramine (H2NF) and a compound NFI were also prepared by this electrolysis. I t has recently been shown that these identifications were incorrect. However, it is to his credit that (in the absence of vapor phase chromatography, mass-, infrared-, nrnr- and epr-spectroscopy), he was able to identify any of the components of the cbmplex gaseous mixture produced by this electrolysis. I n 1942, Haller (3) of Cornell prepared fluorine azide WN3) from the reaction of hydrasoic acid with fluorine. Fluorine azide proved to be an extremely unstable compound which decomposed spontaneously to yield difluorodiazine (NzF2)and nitrogen. Bauer (4) a t Cornell studied the electron diffraction of difluorodiadne and found it to he composed of cis and trans isomers. Tetrafluorohydrazine (NzF4)
Work in the N-F field was quiescent for 11 years until 1958 when Colburn and Kennedy (5) of Rohm and Haas Company, Redstone Research Division, reported the synthesis of tetrafluorohydrazine from the thermal reaction of nitrogen trifluoride with various fluorine acceptors such as copper, bismuth, arsenic, antimony, and stainless steel. The boiling point of tetrafluorohydrazine is -73'C and its critical temperature is 36'C. From an extrapolation of the vapor pressure data a critical pressure of 77 atm is estimated. The infrared absorption spectrum of tetrafluorohydrazine consists of a very strong complex band between 9.75 and 10.75 p and a strong broad band at 13.06 @. Presented as pmt of the Symposium on Recent Advances in Inorganic Chemistry, sponsored jointly by the Divisions of Inorganic Chemistry and Chemical Education, at the 137th Meeting of the ACS, Cleveland, Ohio, April, 1960. This work a~asperformed under the sponsorship of the Army Ordnance Corp. Contract No. 1)A-01-021 ORD-5135.
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The mass spectrum of tetrafluorohydrazine given in Table 1 is consistent with the proposed structure: F
\
N-N
F
/F
F '
The F1# nuclear magnetic resonance spectrum of tetrafluorohydrazine consists of a single broad unresolved band a t a field of approximately 75 ppm lower than that of the Fig nuclei of trifluoroacetic acid. Armstrong (6) and co-workers at the National Bureau of Standards have determined the heat of formation of tetrafluorohydrasine and found it to be -2.0 =t2.5 kcal/mole. Table 1. Fragmentation Pattern of NZFZ M/e
Ion
Pattern
Recently Loughran and Mader (7) of Los Alamos Scientific Laboratory have reportfed an estimated bond dissociation energy of the N-X bond in tetrafluorohydrazine of approximately 30 kcal from appearance potential measurements. Lide and Mann (8) of the National Bureau of Standards have studied the microwave spectrum of tetrafluorohydrazine and report that the observed rotational constants are consistent with a hydrazine-like model (point group C2). Recently several other methods of preparation of tetrafluorohydrazine have been described. Frazer (9) of the Ernest 0. Lawrence Radiation Laboratory has prepared tetrafluorohydrazine from the homogeneous reaction of nitrogen trifluoride with mercury in an electric discharge. Lawton and Weber (10) of Rocketdyne have found that tetrafluorohydrazine is formed when difluoramine (HNF,) is react,ed with various solids (i.e., LiH). Morrow and co-workers (11) at Reaction Motors Division of Thiokol have found that the fluorination of ammonia in a packed reactor produces tetrafluorohydrazinc. Difluoramine (HNF2)
As mentioned above, Ruff (12) reported the formation of difluoramine from the electrolysis of molten ammonium bifluoride. This work has never been confirmed
and indeed, Burg (IS) has pointed out the uncertainty of the identification. Difluoramine was first prepared by Kennedy and Colburn (14) in 1959. The observed trace amounts of difluoramine formed during the preparation of tetrafluorohydrazine by the passage of nitrogen triflnoride (saturated with water vapor) over elemental arsenic at 250-300°C. This preparation of difluoramine suggested the intermediacy of tetrafluorohydrazine and arsine. Further experiments have established the fact that the reaction between arsine and tetrafluorohydrazine does indeed yield difluoramine (15). Difluoramine has also been prepared in yields up to 75% by the reaction of tetrafluorohydrazine with mercaptans (15). Difluoramine has recently been reported arising as a by-product of two fluorination reactions: from the fluorination of urea and subsequent decomposition of intermediate products (If?), and from the fluorination of ammonia (17). Difluoramine is a colorless gas at room temperature. Its boiling point is -23'C and its heat of vaporization is estimated to be 5940 cal/mole. Its melting point is approximately -116'C and its critical temperature (by the Cagniard de la Tour tube method) is 130°C. The infrared absorption spectrum of difluoramine consists of strong doublets a t 7.0, 7.8, 10.2 p and a triplet at 11.2 p. The nmr proton spectrum consists of a triplet, as would be expected from a proton spin-spin coupling with two equivalent fluorine nuclei. The center of the triplet is approximately 6 cycles on the high field side from benzene. The spin-spin splitting is about 24 cycles. The fluorine resonance spectrum consists of two broad bands arising from spin-spin coupling of the fluorine nuclei with the proton. The high field member of this doublet was measured a t 2100 cycles on the low field side of trifluoroacetouitrile. The mass spectrum of difluoramine is given in Table 2. The spectrum was obtained on a Consolidated Electrodynamics Model 620 Spectrometer. Table 2.
Fraqmentation Pattern of HNF,
M/c
Ion
Pattern
(%I
time there is some question as to the identification of the structure of the two isomers which have been isolated. The present discussion assumes that the isolated isomers are cis and trans difluorodiazine.' A third method of preparation of difluorodiazine has been reported by Frazer (E). Diflnorodiazine is produced (along with tetrafluorohydrazine) by the homogeneous reaction of nitrogen trifluoride with mercury in an electric discharge. Morrow et a1 (17), also observed difluorodiazine as a product of the fluorination of ammonia. Low temperature distillation and vapor phase absorption chromatography are t,he best purification methods for difluorodiazine. The trans isomer has been obtained in a purity of 99.7y0 by chromatography and the cis isomer has been obtained by low temperature distillation in a minimum purity of 97.5%. Both physical and chemical properties of the two isomers are quite different as can be seen from the information in Table 3. The mass spectra of the cis and trans isomers of difluorodiazine are showu in Table 4. The FL9nmr spectra of the isomers were observed and are given in Table 5. Table 3.
Physical Properties of Cis and Trans Difluorodiazine
Property
Cis
Tram
Vapor pressure equation
log Pmm = -803.0/T 7.675 -105.7' 3670
log Pmm = -7423/T 7.470 -111.4' 3400
Boiling point ('C) Heat of vaporization (cal/molc) Critical temperature ,on>
i "i
Critical pressure (entimated atm) M~ltingpoint (OC)
+
+
-1"
-13"
70
55
Below - 195"
-172'
The infrared spectrum of the trans isomer in the rock salt region showed a single p-q-r band centering at 10.05 p while the cis isomer showed much more complex absorption with a strong p-q-r band centering at 10.48 p and a strong p-r band at 11.07 and 11.33 p and a very strong gq-r band- centering at 13.57 p. Also weak doublets at 6.52 p and 6.62 p were observed. Table 4.
Mass Spectral Crocking Pattern of Cis and Trons Ditluaradiazine
Dinuorodiazine (N,FI)
Haller's preparation of difluorodiazine (5) from the thermal decomposit,ionof fluorine azide iS not a convenient one because of the danger of explosions during runaway decompositions. Colburn, et al. (18) recently found that difluorodiazine can be prepared by the electrolysis of molten ammonium bifluoride. This observation has been confirmed by Schmeisser (19). Difluorodiazine is a minor product [5-10yo of the condensable (in liquid nitrogen) gases] from the electrolysis of molten (120-13O0C) ammonium bifluoride. As reported by Bauer (4). the difluorodiazine exists in a t least. two isomeric forms, supposedly, eis- and transdifluorodiazine. Fleeting evidence for a third isomeric form of N2Fzhas been reported (18). At the present
66 NPC NJ? 47 33 NF+ 28 Nl 23.5 NxF++ 19 F+ 14 N+ Sensitivity of highest peak div./w +
0.5 100.0 6.0 84.5 1.4 5.3 10.5 12.36
25.3 43.4 5.0 100.0
. ..
1.8 11.6 22.36
A recent report from Lawrence Radiation Laboratory, University of California, Livermore, UCRL-5541 (March 30, 1960), by R. H. Sanborn, presents some spectral evidence concerning the structure of the two isolated isomers. The report confirms the assignment of the trans isomor but supports the 1,l-difluorodiazine structure for the other isolated isomer. The question is still unresolved; hut until it is resolved we shall refer to the second isomeric form as the cw isomer. Volume 38, Number 4, April 1961
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181
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Table 5.
NMR Spectra of Cis and Trans Difluorodiazine
Table 6.
Mass Spectral Crocking Pattern of Chlorarliflumnminc ~
Mle
35)
%j
Ion
Cl
+
Pattern (V!)
21.9
NCl +
6.5 30.1
NF2+
100 26.2 22.4 8.5 1.3 4 6.4 14.5
The chemical reactivities of the cis and trans isomers are quite different. C i s N2F2reacts slowly (completely in two weeks) with glass to form silicon tetrafluoride and nitrous oxide at room temperature and high pressure while the trans isomer is essentially unchanged after a month in a sealed glass tube a t room temperature. Mercury reacts much more rapidly with the &is isomer than with the trans. When heated, the trans isomer is converted into the cis; the heat of isomerization is 27.5 5.0 kcal/mole (18). An extremely interesting reaction of difluorodiazine is its catalvtic effect on nolvmerizat.ion of various monomers. " It has been obkrved that a pressure of 300 mm of N2F2over the monomers of methyl methacrylate, styrene, and cyclopentadiene led to their polymerization in twelve hours or less at room temperature. Diflnorodiazine also catalyzed the polymerization of tetraflnoroethylene at 125°C. At temperatures in the neighborhood of 140°C. polymeric substances were formed in the presence of difluorodiazine with ethylene and propylene. At least in the case of methyl methacrylate, the cis isomer is more effective in catalyzing polymerization than is the trans isomer.
A group of very interesting simple molecules have recently been synthesized and studied: tetrafluorohydrazine (N2F4), difluoramine (HNF*), chlorodifluoramine (NF2C1)and cis and trans diflnorodiazine (N2F2). The further investigation of these compounds should be of great benefit to the understanding of molecular structure and chemical bonding. CAUTION:Numerous explosions have occurred in work with these compounds. Reference is made to the original literature for experimental precautions to be used with these materials.
Chlorodifluoramine (NFXI)
Literature Cited
Petry (80) found in 1960 that when equimolar quantities of difluoramine and boron trichloride are condensed in vacuo a t -13O0C they form a white solid which is at When the 'lid is toward room temperature, decomposition of the solid yields hydrogen chloride (HCI), chlorine (C12), and a new compound, chlorodifluoramine (NF2CI). Chlorodifluoramineis a colorless gas. Its vapor pressure curve is given by the equation log Pmm =
-+ 7.478.
-950
The extrapolated boiling point is -67'C. The heat of vaporization calculated from the above equation is 4350 cal/mole. The melting point of chlorodiflnoramine has not been obtained but lies between -183' and -196°C. The mass spectrum of chlorodifluoramine is given in Table 6. The infrared spectrum of chlorodifluoramine consists of very strong hands centered a t 10.8 (triplet), 11.7 (doublet), and 14.4 p (triplet); a doublet of moderate intensity centered a t 13.4 p and weak hands at 5.4, 5.7, 5.9, and 7.3 p. The FlS nmr s ~ e c t r u mof chlorodifluoramine consists of a single broad hand centered a t 8685 cycles to the low field side of trifluor~acet~ic acid. Uses of N-F
Compounds
The o d y use of N-F compounds reported in the literature is the use of nitrogen trifluoride with hydrogen in a torch (81). This torch can weld, braze, and cut 182
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51 52
%}
3 89
BFClf GIl+ NFzClt
flux since t,he nitrogen tri. metals without fluoride acts as a gaseous flux. Summary
(1) RUFF,O., FISCFIER,J., AND LUFT,F., Z. Anory. Allgem. Chem., 172, 417 (1928). (2) Rum, O., AND GEISEL,E., Ber., 36, 2677 (1903). J. F., Ph.1). Thesis, Cornell University, 1942. (3) HALLER, (4) BAUER,8 . H., J. Am. C h m . Soe., 69,3104 (1947). (5) COLBURN, C. B., AND KENNEDY, A., J. Am. Chem. Soe., 80, 5004 (1958). 1 (6) ARMSTRONG, G. T.,MARANTZ, S., AND COYLE,C. F., National Bureau of Standards Report No. 6584, October, a,
10co -7-7.
(7) LOUOHRAN, E. D., AND MADER,C., J. Chem. Phya., 32, 1578 (1960). (8) LIDE, D. R., JR., AND MANN,D. E., J . Chem. Phys., 31, 1129 (1959). (9) FRAZER, J. W., J. I ~ O Tand B Nue. Chm., 11, 168 (1959). (10) LAWTON,E. A.2 AND WEBE%J. Q., J . Am. Chem. So& 81, 4755 (1959). (11) MORROW, S. I., PERRY,D. D., A ~ DCOHEN,M., J . Am. Chem. Soc., 81, 6338 (1959). (12) RUFF, O., AND STAUB,L., Z. Anorg. Allgem. Chem., 198, 32 (1931). (13) BURG, A. B., i n Simons, "Fluorine Chemistry," Val. 1, Academic Press, New York, 1950, p. 88. (14) KENNEDY, A,, AND COLBURN, C. B., J . Am. Chmt. Soe., 81, 2906 (1959). (15) FREEMAN,J. P., KENNEDY,A.2 AND COL~URN~ C. B.2 J . Am. Chem. Soc., (m press). (16) LAWTON, E, A,, AND J, Q,, J , Am. Chm. See., 81, 4i55 (1959). 117) S. I.. PERRY.D. D.. COHEN.M. S.. AND SCHOEN. . MORROW. FELDE$,C. 'w., ~ b s t r a c t s 137th , Meeting of the ACS, Cleveland, Ohio, April, 1960, p. 11M. C.B., JOHNSON, F. A,, KENNEDY, A., MCCALLUM, (18) COLBURN, K. S., METZGER,L. C., A N D PARKER,C. O., J. Am. C h m . Soe., 81, 6397 (1959). M., AND SARTORI, P., Angem. Chem., 71, 523 (19) SCHMEISSER, (1959). (20) PETRY,R. C., J. Am. Chem. Soe., 82, 2400 (1960). H. H., Ind. and Eng. Chem., 51, 309 (1959) (21) ROGERS,
