15 The Chemistry of Difluoramines A. D. C R A I G , G. A. W A R D , C. M. W R I G H T , and J. C. W. C H I E N
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Research Center, Hercules Powder Co., Wilmington, Del.
The chemistry of difluoramines of the type X—NF has been studied to gain an understanding of the nature of the N — F and N—X bonds, to obtain a picture of the relative electron distributions in X—NF compounds, and to determine the existence and stabilities of N—F radicals and ions. These compounds have been studied using electrochemistry, complexation, infrared spectroscopy, and theoretical calculations. Oxidation-reduction reactions have been carried out, and the effects of various environments on the N—F and N—X bonds have been investigated. The results of these studies emphasize the chemistry of difluoramine and the existence and 2
2
NF ,
stability of
+
2
·NF , 2
NF ,
H NF .
and
-
2
+
2
2
/Compounds containing the N F group are commonly referred to as ^ difluoramines. The chemistry of these compounds has been studied more intensively during recent years than previously and has been the subject of three review articles (3, 9,14). The aims of our work, part of which we describe here, are to obtain a picture of the relative electron distributions i n X — N F compounds, to gain an understanding of the n a ture of the N — F and N — X bonds, and to determine the existence and stabilities of N — F radicals and ions. W e have concentrated on X — N F compounds, where X = F , C l , H , N F , C H , C H , and C F . The radical and ionic species which have held our attention are N F , N F , N F + , N F " ~ , N F + , N F ~ , N F + , and H N F + . This paper emphasizes the chemis try of difluoramine, H N F , and the existence of NF -containing ions to illustrate our investigations i n N F chemistry. 2
2
2
2
3
2
5
3
e
2
2
3
2
e
2
2
2
2
2
Results and
Discussion
Theoretical Considerations. One approach we have taken is calcu lating molecular parameters of various N — F species b y using molecular orbital treatments (2,7). The π-bond orders and atomic charges calcuHolzmann; Advanced Propellant Chemistry Advances in Chemistry; American Chemical Society: Washington, DC, 1966.
15.
CRAIG ET AL.
1 49
Chemistry of Difluoramines
lated for N F moieties (Table I ) are particularly relevant to N F chem istry. 2
2
Table I.
Calculated Charges and Bond Orders of NF Species 2
Atomic Charges
w-Bond
Q -0.353 -0.43 -0.5
Orders 0.3 0.35 0
F
NF + NF NF ~
+1.706 +0.86 0
2
2
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2
Kaufman and co-workers have made similar calculations and have shown that i n covalent N F compounds there are no orbitals available on the nitrogen having energies low enough for significant ττ-bonding with the unshared electrons on the fluorine atoms (12, 13). I n a series of X - N F compounds, the relative electron distributions are thus a function of the inductive effects of the X group. The various situations are envisioned as follows: 2
2
p-0J53
+ 1.706 Ν
/
ρ-0.43
+0.86 Ν
\
ρ-0.5
/
F-0-JM
0.0 Ν
\
/
\
F-0-43
I
F-0-6
II
III
F
^
krs /
F5-
/~% ( /
Χ—Ν
Χ—Ν
F
^> F5_
IV
V
The relatively high negative charge on fluorines causes many N F species to decompose readily by losing a fluoride ion. This work indicates that the N F - c o n t a i n i n g ions which have the greatest probability of long term existence are the cations because N F + , of all the N F species, has the lowest negative atomic charge on the fluorines. I n all of the difluor amine derivatives studied, the nonbonding L-shell electrons of the nitrogen are more displaced toward the fluorines than strictly localized on the nitrogen. As a result, the nitrogen i n many N F species carries a relatively high positive charge. A n y cationic N F species would be expected to have a high electron affinity, probably higher than that of N 0 + . In isolating an N F cation, one must stabilize it w i t h a large anion of low charge density so that polarization can occur or one must select a hypothetical N F salt i n which the lattice stabilization energy is very high. Candidates for the latter are moderately sized divalent anions. N o NF -containing ions of any type have been observed except for N F + which has been observed i n the mass spectrometer (8). 2
2
2
2
2
2
2
2
2
2
2
Holzmann; Advanced Propellant Chemistry Advances in Chemistry; American Chemical Society: Washington, DC, 1966.
1 50
ADVANCED PROPELLANT CHEMISTRY
Experimental Observations. T h e nearest we have come to demon strating the existence of N — F ions is i n the electrochemical oxidationreduction reactions of H N F . The oxidation has been carried out i n water and i n various polar organic solvents under acid conditions (15). T h e first step of this reaction is formation of the * N F radical. T h e * N F radical undergoes combination processes on the surface of the electrode rather than diffusing into the body of the solution before being involved in further reactions. The combination process on the electrode surface has been used to prepare various N F compounds by simultaneously generat ing other radical species—e.g., 2
2
2
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2
HNF
H
2
GHjCOOH •NFj +
+
+ e~ +
NF
CH, + H
-CH,
CHfNF
2
+
+ C O , + e"
2
The oxidation of H N F involves the removal of an electron from the nitrogen i n a solvated H N F species rather than from N F ~ . A l l of our work on the solution chemistry of H N F show that under conditions favor ing the formation of N F " ~ (proton removal), this species loses a fluoride ion to form difluorodiazine. 2
2
2
2
2
NF " 2
7 N F 2
2
2
+ F "
The reduction of H N F i n aqueous media is a four-electron process i n which H N F is reduced to ammonia (16). I n nonaqueous solvents the reduction depends on the availability of protons i n the system, and the reduction potential is strongly influenced by the degree and type of solva tion ( Table II ). Nitrogen trifluoride is electrolytically reduced i n aqueous solution at —1.40 volts vs. S.C.E. Six electrons per molecule of N F are involved i n the reduction. 2
2
3
HNF
2
+ 4e~ + 2H+
N F , + 6e- + 3H+
— —
NH, + 2F" N H , + 3F~
N F could not be electrolytically oxidized i n systems similar to those used for the reduction. 3
Table II· Effect of Solvent on the Polarographic Reduction of HNF £i/2
Solvent
2
HNFt
(volts vs. S.C.E.)
H 0 CH OH GHiCN Dimethylformamidc Dimethyl sulfoxide 2
8
1.22 1.55 1.42 1.61 1.64
The solvation of H N F i n various solvents was studied b y conven tional and low temperature infrared techniques and b y determining the dissociation pressure-temperature relationships of several H N F solvent complexes (16). Total enthalpies of dissociation were determined where 2
2
Holzmann; Advanced Propellant Chemistry Advances in Chemistry; American Chemical Society: Washington, DC, 1966.
15.
CRAIG ET AL.
Chemistry of Difluoramines
1 51
experimentally feasible. Shifts i n the Ν—H and N — F stretching fre quencies i n the I R spectra of 1:1 complexes of H N F with solvents and of lM-solutions of H N F were examined to determine the nature of bonding i n the solvated species (Table I I I ) . I n general these data indicate the order H 0 < C H O H < C H C N < H C O N H < H O C N ( C H ) ( C H ) S O for the strength of solvation of H N F . It was found that 1:2 complexes of H N F with D M F and D M S O exert very little vapor pressure at room temperature. Equimolar complexes of H N F with formamide, dimethylformamide, and dimethyl sulfoxide exert relatively little vapor pressure at 0° C . H N F is b y far the least solvated b y water. 2
2
2
3
3
8
2
2
8
2
2
2
2
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2
Table III.
Infrared Absorption Maxima of HNF Complexes 2
N—H Stretch
Material C H , C N H N F i (solid)* \M H N F , in G H i O H (liquid)* H j O · H N F s (solid)* \M H N F in D M S O (liquid)* H N F j (solid)* 2
N- - F Asym. Sym. Stretch Stretch
e
1325» 1320
e
960 955 973
860 855 875
6
β
952
850 855 880
1424
2710 2725 2800 2975 2725 2740 3110
D M F - H N F Î (solid)*
N- -H Sym. Asym. Bend Bend
1390
e 1450
e
1350
β
972
a A t —160° C At 2 5 ° C . * Solvent interference
h
In a l l of the solvents studied, the most important factor, determined from vibrational spectra, is the strength of the difluoramine hydrogen bond with the solvent. Thus, H N F appears to be bonded to these solvents i n structures of the type 2
F CHjCN
/
Η—Ν
\
F
However, i n some systems the infrared spectra indicate moderate bonding of the fluorines with the solvent—e.g., H,C
O....H ^N—
