Synthesis and characterization of salts containing the BrO 3F2- anion; a rare example of a bromine(VII) species

Article abstract of DOI:10.1021/ja0402607

The BrO₃F₂^- anion has been prepared by reaction of BrO₃F with the fluoride ion donors KF, RbF, CsF, [N(CH₃)₄][F], and NOF. The BrO₃F ₂^- anion is only the fourth Br(VII) species to have been isolated in macroscopic quantities, and it is one of only three oxide fluorides that possess D_3h symmetry, the others being XeO₃F ₂ and OsO₃F₂. The fluoride ion acceptor properties of BrO₃F contrast with those of ClO₃F, which does not react with the strong fluoride ion donor [N(CH₃) ₄][F] to form the analogous ClO₃F₂^- salt. The single-crystal X-ray structures of [NO]₂[BrO ₃F₂][F] and [N(CH₃)₄][BrO ₃F₂] confirm the D_3h, symmetry of the BrO ₃F₂^- anion and provide accurate Br-O (1.593(3)-1.610(6) A) and Br-F (1.849(5)-1.827(4) A) bond lengths. The salt, [NO]₂[BrO₃F₂][F], is fully ordered, crystallizing in the monoclinic space group, C2/c, with a = 9.892(3) A, b = 12.862(4) A, c = 10.141(4) A, β = 90.75(2)°, V = 12460(7) A³, Z = 4, and R₁ = 0.0671 at -173 °C, whereas [N(CH₃)₄)][BrO₃F₂] exhibits a 2-fold disorder of the anion, crystallizing in the tetragonal space group, P4/nmm, with a = 8.5718(7) A, c = 5.8117(6) A, V = 427.02(7) A³, Z = 2, and R₁ = 0.0314 at -173 °C. The ¹⁹F chemical shift of [N(CH₃)₄)][BrO ₃F₂] in CH₃CN is 237.0 ppm and is more deshielded than those of the previously investigated Br(VII) species, BrO ₃F and BrF₆⁺. The vibrational frequencies of the BrO₃F₂^- anion were determined by use of Raman and infrared spectroscopy and were assigned with the aid of electronic structure calculations and by analogy with the vibrational assignments reported for XeO₃F₂ and OsO₃F₂. The internal and symmetry force constants of BrO₃F₂^- were determined by use of general valence force field and B-matrix methods, respectively, and are compared with those of XeO₃F₂, OsO₃F₂, and the unknown ClO₃F₂ ^- anion. The instability of ClO₃F₂^- relative to BrO₃F₂^- has been investigated by electronic structure calculations and rationalized in terms of atomic charges, Mayer bond orders, and Mayer valencies, and the enthalpies of fluoride ion attachment to BrO₃F and ClO₃F.

Full text of DOI:10.1021/ja0402607

Published on Web 06/08/2005
Synthesis and Characterization of Salts Containing the
-
BrO3F2 Anion; A Rare Example of a Bromine(VII) Species
John F. Lehmann and Gary J. Schrobilgen*
Contribution from the Department of Chemistry, McMaster UniVersity,
Hamilton, Ontario, L8S 4M1, Canada
Received November 23, 2004; E-mail: schrobil@mcmaster.ca
-
Abstract: The BrO3F2 anion has been prepared by reaction of BrO3F with the fluoride ion donors KF,
RbF, CsF, [N(CH3)4][F], and NOF. The BrO3F2- anion is only the fourth Br(VII) species to have been isolated
in macroscopic quantities, and it is one of only three oxide fluorides that possess D3h symmetry, the others
being XeO3F2 and OsO3F2. The fluoride ion acceptor properties of BrO3F contrast with those of ClO3F,
-
which does not react with the strong fluoride ion donor [N(CH3)4][F] to form the analogous ClO3F2 salt.
The single-crystal X-ray structures of [NO]2[BrO3F2][F] and [N(CH3)4][BrO3F2] confirm the D3h symmetry of
-
the BrO3F2 anion and provide accurate Br-O (1.593(3)-1.610(6) Å) and Br-F (1.849(5)-1.827(4) Å)
bond lengths. The salt, [NO]2[BrO3F2][F], is fully ordered, crystallizing in the monoclinic space group, C2/c,
with a ) 9.892(3) Å, b ) 12.862(4) Å, c ) 10.141(4) Å, â ) 90.75(2)°, V ) 12460(7) Å3, Z ) 4, and R1 )
0.0671 at -173 °C, whereas [N(CH3)4)][BrO3F2] exhibits a 2-fold disorder of the anion, crystallizing in the
tetragonal space group, P4/nmm, with a ) 8.5718(7) Å, c ) 5.8117(6) Å, V ) 427.02(7) Å3, Z ) 2, and R1
) 0.0314 at -173 °C. The 19F chemical shift of [N(CH3)4)][BrO3F2] in CH3CN is 237.0 ppm and is more
deshielded than those of the previously investigated Br(VII) species, BrO3F and BrF6+. The vibrational
frequencies of the BrO3F2- anion were determined by use of Raman and infrared spectroscopy and were
assigned with the aid of electronic structure calculations and by analogy with the vibrational assignments
reported for XeO3F2 and OsO3F2. The internal and symmetry force constants of BrO3F2- were determined
by use of general valence force field and B-matrix methods, respectively, and are compared with those of
XeO3F2, OsO3F2, and the unknown ClO3F2- anion. The instability of ClO3F2- relative to BrO3F2- has been
investigated by electronic structure calculations and rationalized in terms of atomic charges, Mayer bond
orders, and Mayer valencies, and the enthalpies of fluoride ion attachment to BrO3F and ClO3F.
BrO2+16,17), and BrOF3 (BrOF4- 18,19 BrOF2+20,21). Fluoride ion
,
Introduction
transfer reactions, however, have not been used to prepare new
cations and anions derived from known Br(VII) species. Among
Br(VII) species, BrF6+ and BrO3F might be expected to accept
fluoride ions leading to BrF7 and BrO3F2- or to act as a fluoride
ion donor leading to BrO3+. Although a high-temperature
synthesis of BrF7 has been reported,22 the existence of BrF7
Prior to the present study, the only Br(VII) species that had
been prepared in macroscopic quantities and characterized were
6,7
BrO3F,1 salts of the BrF6+ 2,3 and BrO4-4-7 (including HBrO4
, )
ions.8 The synthetic challenges associated with the syntheses
of Br(VII) compounds are consistent with the general reluctance
of late row four elements of the periodic table to attain their
highest oxidation states.
+
was shown to be unlikely because the reaction of BrF6 with
NOF at -78 °C results in reduction to BrF6- with the evolution
of F2,3 and because attempts to form BrF7 by oxidation of BrF5
with F2 under photolytic conditions at -40 to -60 °C have
also failed.23 The fluoride ion donor properties of BrO3F have
The ability of main-group and transition-metal fluorides and
oxide fluorides to accept or donate fluoride ions is well-known,
providing routes to anionic and cationic derivatives of BrF3
(BrF2+,9,10 BrF4-11), BrF5 (BrF6-,12 BrF4+13,14), BrO2F (BrO2F2-,15
(12) Bougon, R.; Charpin, P.; Soriano, J. C. R. Hebd. Se´ances Acad. Sci. Ser.
C 1971, 272, 565.
(1) Appelman, E. H. Studier, M. H. J. Am. Chem. Soc. 1969, 91, 4561.
(2) Gillespie, R. J.; Schrobilgen, G. J. J. Chem. Soc., Chem. Commun. 1974,
90.
(3) Gillespie, R. J.; Schrobilgen, G. J. Inorg. Chem. 1974, 13, 1230.
(4) Brown, L. C.; Begun, G. M.; Boyd, G. E. J. Am. Chem. Soc. 1969, 91,
2250.
(5) Appelman, E. H. J. Am. Chem. Soc. 1968, 90, 1900.
(6) Appelman, E. H. Inorg. Synth. 1971, 13, 1.
(7) Appelman, E. H. Inorg. Chem. 1969, 8, 223.
(8) Isopropyl perbromate, (CH3)2CHOBrO3, has been reported but is poorly
characterized (Baum, K.; Beard, C. D.; Grakauskas, V. J. Am. Chem. Soc.
1975, 97, 267).
(9) Woolf, A. A.; Emele´us, H. J. J. Chem. Soc. 1949, 2865.
(10) Brown, D. H.; Dixon, K. R.; Sharp, D. W. A. Chem. Commun. 1966, 654.
(11) Siegel, S. Acta Crystallogr. 1956, 9, 493.
(13) Lind, M. D.; Christe, K. O. Inorg. Chem. 1972, 11, 608.
(14) Vij, A.; Tham, F. S.; Vij, V.; Wilson, W. W.; Christe, K. O. Inorg. Chem.
2002, 41, 6397.
(15) Gillespie, R. J.; Spekkens, P. H. J. Chem. Soc., Dalton Trans. 1977, 1539.
(16) Jacob, E. Angew. Chem., Int. Ed. Engl. 1976, 15, 158; Angew. Chem. 1976,
88, 189.
(17) Spekkens, P. H. Ph.D. Thesis, McMaster University, Hamilton, Ontario,
Canada, 1977; pp 87-104.
(18) Gillespie, R. J.; Spekkens, P. J. Chem. Soc., Dalton Trans. 1976, 2391.
(19) Bougon, R.; Joubert, P.; Tantot, G. J. Chem. Phys. 1977, 66, 1562.
(20) Adelhelm, E.; Jacob, E. Angew. Chem., Int. Ed. Engl. 1977, 16, 461; Angew.
Chem. 1977, 89, 476.
(21) Bougon, R.; Huy, T. B.; Charpin, P.; Gillespie, R. J.; Spekkens, P. H. J.
Chem. Soc., Dalton Trans. 1979, 6.
(22) Fogle, C. E.; Rewick, R. T. U.S. Patent 3615206, 1971.
9
9416
J. AM. CHEM. SOC. 2005, 127, 9416-9427
10.1021/ja0402607 CCC: $30.25 © 2005 American Chemical Society

-
Synthesis of Salts Containing the BrO3F2 Anion
A R T I C L E S
been assessed in anhydrous HF solvent by monitoring solutions
of BrO3F and SbF5 by Raman spectroscopy, and solutions of
BrO3F and AsF5 by 19F NMR spectroscopy.24 The absence of
vibrational frequency shifts or changes in the 19F NMR chemical
shift of BrO3F in these studies indicates that interactions between
the ligand atoms of BrO3F and the Lewis acids are weak to
negligible. Although the fluoride ion acceptor properties of
BrO3F have not been explicitly investigated, the observation
that CsF and NO2F do not behave as fluoride ion donors toward
ClO3F25 has led to speculation that BrO3F may also be a poor
fluoride ion acceptor.24 Recent theoretical calculations,26 how-
ever, suggest that BrO3F may react with strong fluoride ion
donors such as [N(CH3)4][F], so-called “naked fluoride”.27
The VSEPR model of molecular geometry28 predicts that the
BrO3F2 and ClO3F2 anions should possess trigonal bipyr-
amidal geometries (D3h symmetry) in which the three oxygen
atoms occupy the equatorial positions and the two fluorine atoms
occupy the axial positions. Although the trigonal bipyramidal
geometry is common among neutral, gas-phase, and anionic
inorganic pentahalides,29 the only trigonal bipyramidal EO3F2
species known are XeO3F230 and OsO3F2.31,32 Moreover, OsO3F2
is a polymeric cis-fluorine bridged chain structure in the solid
state,33 but has been shown by vibrational spectroscopy to exist
as a monomer having D3h symmetry when isolated in a
matrix.31,32 The challenging synthesis of XeO3F2 by the reaction
of XeO4 with XeF6 has yet to yield bulk quantities of pure
XeO3F2, but 19F NMR spectroscopy of SO2ClF, HF, and BrF5
solutions of XeO3F2,34 and vibrational (Raman, IR) spectra of
XeO3F2 isolated in Ar and Ne matrices30 have confirmed the
predicted D3h geometry.35 The geometric parameters of XeO3F2
and the OsO3F2 monomer have not been determined.
cally less favorable than fluoride ion transfer to HF. Conse-
quently, the use of HF as a solvent was abandoned for the
remainder of these studies.
Although violent detonations occasionally occurred when
BrO3F was brought into contact with CH3CN, the high first
adiabatic ionization potential (13.59 eV),36 favorable liquid
range, and resistance to fluoride ion attack at low tempera-
tures37,38 make CH3CN a useful solvent for the synthesis of
anions derived from strong oxidative fluorinators.39,40 The
N(CH3)4+, K+, Rb+, and Cs+ salts of BrO3F2- were synthesized
(eq 1) and handled in CH3CN solvent for prolonged periods of
time at low temperatures (<-30 °C) and at ambient tempera-
tures for short periods of time.
-
-
MF + BrO3F f [M][BrO3F2] (M ) K, Rb, Cs, N(CH3)4)
(1)
Raman spectroscopy was used to verify salt formation prior to
solvent removal under dynamic vacuum at -40 °C. The
reactivities of [N(CH3)4][F] and CsF toward BrO3F were similar,
with [N(CH3)4][BrO3F2] and [Cs][BrO3F2] formation proceeding
smoothly over a period of an hour at -40 to -48 °C. The low-
temperature R-phase of [Cs][BrO3F2] was isolated when the
solution was maintained at or below -40 °C during reaction
and removal of the solvent, whereas the high-temperature
â-phase of [Cs][BrO3F2] was isolated when the reaction and
solvent removal were carried out at -35 °C. The â-phase of
[Cs][BrO3F2] may alternatively be obtained by allowing
R-[Cs][BrO3F2] to stand at 0 °C for 35 h in the absence of a
solvent; however, the reverse process does not occur at
temperatures as low as -78 °C over a period of 4 days.
The reactions of KF and RbF with BrO3F were slow at -40 °C
when compared with the nearly instantaneous formation of
[N(CH3)4][BrO3F2] and R-[Cs][BrO3F2] at or below this tem-
perature, but proceeded rapidly at -30 to -35 °C. The slower
reaction rates of KF and RbF at temperatures below -40 °C
likely reflect their lower solubilities.
The present work investigates the fluoride ion acceptor
behavior of BrO3F toward a range of fluoride ion donors, with
the view to synthesize salts of the BrO3F2- anion, thus providing
a significant extension of Br(VII) chemistry and the first
example of a trioxide difluoride anion. The study also reinves-
tigates the fluoride ion acceptor properties of ClO3F.
The double salt, [NO]2[BrO3F2][F], was prepared by the
reaction of BrO3F with neat NOF at -78 °C (eq 2) for 5 min
followed by rapid removal (ca. 30 s) of excess NOF under
dynamic vacuum at -78 °C.
Results and Discussion
Syntheses of [M][BrO3F2] (M ) K, Rb, Cs, N(CH3)4) and
[NO]2[BrO3F2][F]. The fluoride ion acceptor properties of
BrO3F were investigated by monitoring its reactions with NOF
and MF by means of Raman spectroscopy. An initial attempt
to synthesize [Cs][BrO3F2] in anhydrous HF failed to provide
Raman spectroscopic evidence for fluoride ion transfer to BrO3F,
suggesting that fluoride ion transfer to BrO3F is thermodynami-
2NOF + BrO3F f [NO]2[BrO3F2][F]
(2)
The colorless product had a significant dissociation vapor
pressure at -78 °C and was crystallized by allowing the solid
to sublime under 1000 Torr of Ar over a period of several weeks
at -78 °C. The Raman spectra of the initial and crystallized
products were identical, and single-crystal X-ray diffraction was
used to establish the product stoichiometry in eq 2.
Attempted Syntheses of [N(CH3)4]2[BrO3F3] and [N(CH3)4]-
[ClO3F2]. The reaction of [N(CH3)4)][BrO3F2] with a stoichio-
metric amount of [N(CH3)4)][F] was attempted at 0 °C in
CH3CN solvent. Upon removal of the solvent at -40 °C, the
Raman spectrum of the remaining solid showed that it was a
(23) Pilipovich, D.; Rogers, H. H.; Wilson, R. D. Inorg. Chem. 1972, 11, 2192.
(24) Gillespie, R. J.; Spekkens, P. H. Isr. J. Chem. 1978, 17, 11.
(25) Christe, K. O.; Schack, C. J. AdV. Inorg. Chem. Radiochem. 1976, 18, 319.
(26) Christe, K. O.; Dixon, D. A. Presented at the 16th Winter Fluorine
Conference, St. Petersburg, FL, Jan 12-17, 2003.
(27) Christe, K. O.; Jenkins, H. D. B. J. Am. Chem. Soc. 2003, 125, 9457.
(28) Gillespie, R. J.; Hargittai, I. In The VSEPR Model of Molecular Geometry;
Allyn and Bacon: Boston, 1991.
(29) Nakamoto, K. Infrared and Raman Spectra of Inorganic and Coordination
Compounds, Part A, 5th ed.; Wiley & Sons: New York, 1997; pp 209-
211.
(30) Claassen, H. H.; Huston, J. L. J. Chem. Phys. 1971, 55, 1505.
(31) Beattie, I. R.; Blayden, H. E.; Crocombe, R. A.; Jones, P. J.; Ogden, J. S.
J. Raman Spectrosc. 1976, 4, 313.
(32) Hope, E. G.; Levason, W.; Ogden, J. S. J. Chem. Soc., Dalton Trans. 1988,
61.
(36) Dibeler, V. H.; Liston, S. K. J. Chem. Phys. 1968, 48, 4765.
(37) Christe, K. O.; Wilson, W. W. J. Fluorine Chem. 1990, 47, 117.
(38) Christe, K. O.; Wilson, W. W.; Wilson, R. D.; Bau, R.; Feng, J. J. Am.
Chem. Soc. 1990, 112, 7619.
(39) Emara, A. A. A.; Schrobilgen, G. J. Inorg. Chem. 1992, 31, 1323.
(40) Schrobilgen, G. J. J. Chem. Soc., Chem. Commun. 1988, 863.
(33) Bougon, R.; Buu, B.; Seppelt, K. Chem. Ber. 1993, 126, 1331.
(34) Gerken, M.; Schrobilgen, G. J. Coord. Chem. ReV. 2000, 197, 335.
(35) Gillespie, R. J. In Noble-Gas Compounds; Hyman, H. H., Ed.; University
of Chicago Press: Chicago, 1963; pp 333-339.
9
J. AM. CHEM. SOC. VOL. 127, NO. 26, 2005 9417

A R T I C L E S
Lehmann and Schrobilgen
Table 1. Summary of Crystal Data and Refinement Results for
[NO]2[BrO3F2][F] and [N(CH3)4][BrO3F2]
Table 2. Bond Lengths and Bond Angles for [NO]2[BrO3F2][F] and
[N(CH3)4][BrO3F2]
[NO]2[BrO3F2][F]
[N(CH3)4][BrO3F2]
[NO]2[BrO3F2][F]
space group
a (Å)
b (Å)
C2/c
P4/nmm
8.5718(7)
8.5718(7)
5.8117(6)
90
427.02(7)
2
240.03
1.867
-173
4.81
0.0314
0.0761
bond lengths (Å)
bond angles (deg)
9.892(3)
12.862(4)
10.141(4)
105.38(2)
1246.0(7)
4
244.88
2.616
-173
6.66
Br(1)-O(1)
1.599(7)
1.602(5)
1.872(4)
1.593(7)
1.610(6)
1.849(5)
1.038(9)
1.066(8)
2.746
2.854
2.852
2.631
2.847
3.000
2.544
2.827
2.032
2.216
2.412
2.214
2.672
O(1)-Br(1)-O(2)
119.7(2)
90.4(1)
179.3(3)
89.9(3)
118.2(3)
90.8(2)
89.8(3)
178.5(3)
132.1
143.8
90.8
136.3
116.3
c (Å)
Br(1)-O(2)
Br(1)-F(1)
Br(2)-O(3)
Br(2)-O(4)
Br(2)-F(2)
N(1)-O(5)
O(1)-Br(1)-F(1)
F(1)-Br(1)-F(1A)
O(2)-Br(1)-F(1)
O(3)-Br(2)-O(4)
O(3)-Br(2)-F(2)
O(4)-Br(2)-F(2)
F(2)-Br(2)-F(2A)
N(1A)-O(5A)‚‚‚F(1)
N(2B)-O(6A)‚‚‚F(1)
O(6B)-N(2A)‚‚‚F(1)
N(1B)-O(5B)‚‚‚F(2)
O(6A)-N(2A)‚‚‚F(2)
O(5A)-N(1A)‚‚‚F(2)
O(5)-N(1)‚‚‚F(3)
â (deg)
V (Å3)
Z
mol wt (g mol-1
)
F
calcd (g cm-3
)
T (°C)
N(2)-O(6)
µ (mm-1
R1
wR2
)
O(5A)‚‚‚F(1)
O(6A)‚‚‚F(1)
O(6B)‚‚‚F(1)
N(2A)‚‚‚F(1)
O(5B)‚‚‚F(2)
O(5A)‚‚‚F(2)
N(1A)‚‚‚F(2)
N(2B)‚‚‚F(2)
N(1)‚‚‚F(3)
N(2)‚‚‚F(3)
N(1A)‚‚‚F(3)
N(2A)‚‚‚F(3)
O(6)‚‚‚F(3)
O(6A)‚‚‚F(3)
O(5)‚‚‚F(3)
O(5A)‚‚‚F(3)
a
0.0671
0.1813
b
a R1 ) ∑||Fo| - |Fc||/∑(|Fo|) for I > 2σ(I). b wR2 ) ∑|(|Fo| - |Fc|)w1/2|/
∑(|Fo|w) for I > 2σ(I).
105.9
105.5
114.4
103.4
O(5A)-N(1A)‚‚‚F(3)
O(6)-N(2)‚‚‚F(3)
O(6A)-N(2A)‚‚‚F(3)
mixture of [N(CH3)4][BrO3F2] and [N(CH3)4)][F], providing no
evidence for [N(CH3)4]2[BrO3F3] formation.
102.3
The synthesis of [N(CH3)4][ClO3F2] was attempted by the
reaction of [N(CH3)4][F] at -40 °C with neat ClO3F and with
ClO3F dissolved in CH3CN solvent. Raman spectroscopy failed
to provide evidence for [N(CH3)4][ClO3F2] formation, which
is consistent with earlier attempts to synthesize salts of the
2.654
2.517
2.954
-
ClO3F2 anion by reaction of ClO3F with the weaker fluoride
[N(CH3)4][BrO3F2]
ion donors CsF25 and NO2F.25
bond lengths (Å)
bond angles (deg)
X-ray Crystal Structures of [NO]2[BrO3F2][F] and
[N(CH3)4][BrO3F2]. Unit cell parameters and refinement sta-
tistics for [NO]2[BrO3F2][F] and [N(CH3)4][BrO3F2] at -173
°C are given in Table 1, and geometric parameters are given in
Table 2.
Br-O(1)
Br-O(2)
Br-F(1)
N-C
1.600(3)
O(1)-Br-O(2)
O(1)-Br-F(1)
O(2)-Br-F(1)
C-N-C(A)
118.7(5)
98.5(3)
85.9(4)
109.2(1)
110.0(2)
1.44(1)
1.77(6)
1.496(2)
0.924
0.908
4.21
C-H(1)
C-N-C(B)
C-H(2)
BrO(1)‚‚‚Br(B)
(a) [NO]2[BrO3F2][F]. The X-ray crystal structure of
[NO]2[BrO3F2][F] (Figure 1) is fully ordered, providing accurate
-
geometric parameters for the BrO3F2 anion. The asymmetric
unit of this salt is best described in terms of two crystallo-
respectively. The remaining contacts to F(1) involve both atoms
of the same NO+ cation with N‚‚‚F and O‚‚‚F distances of 2.631
and 2.852 Å, respectively, and an O-N‚‚‚F angle of 90.8°. Two
contacts to F(2) involve both atoms of the same NO+ cation,
with N‚‚‚F and O‚‚‚F distances of 2.544 and 3.000 Å,
respectively, and an O-N‚‚‚F angle of 105.9°. The two
remaining contacts to F(2) are of the N-O‚‚‚F (2.847 Å) and
O-N‚‚‚F (2.827 Å) types with N-O‚‚‚F and O-N‚‚‚F angles
of 136.3 and 116.3°, respectively. The O‚‚‚F contacts to both
crystallographically independent anions are close to the sum of
the fluorine and oxygen van der Waals radii (2.75,41 2.99 Å42)
and are expected to be long as a result of the negative charges
on the oxygen and fluorine atoms. In contrast, the N‚‚‚F contacts
are shorter than the sum of the fluorine and nitrogen van der
Waals radii (2.85,41 3.08 Å42). Although there are two N‚‚‚F
contacts to F(2), with the shorter of these (2.544 Å) being nearly
0.10 Å less than the single N‚‚‚F(1) (2.631 Å) contact, the Br-
F(1) bond is significantly longer than the Br-F(2) bond. The
difference may arise from the π-acceptor properties of NO+
that result from its σ2gσu/2σg2πu4πg/0 electronic configuration.
Transfer of electron density from the anion to the cation is
expected when the LUMO (π/g0) molecular orbital is directed
toward the electron donor atom. The ideal π-acceptor orientation
for NO+ only occurs for F(1) (O(6B)-N(2A)‚‚‚F(1), 90.8°) and
-
graphically unique BrO3F2 and NO+ ions and a single F-
anion, with each NO+ and F- ion having a symmetry-related
position.
-
The BrO3F2 anion geometry is based upon a trigonal
bipyramidal VSEPR arrangement in which the larger steric
demands of the double bond pair domains of the oxygen ligands
require that they occupy the equatorial positions. The minor
distortions of the O-Br-O (118.2(2)-123.5(5)°) and O-Br-F
(89.4(3)-90.7(2)°) bond angles from their ideal 120 and 90°
values are attributed to solid-state packing effects. The position
of the Br-O(3) bond on the 2-fold axis results in each anion
having two unique Br-O bonds and two equivalent Br-F
bonds. As a result, there are four unique Br-O bond lengths
that are the same within (3σ, with a range of 1.593(7)-1.610(5)
Å and an average value of 1.601(7) Å.
The Br-F bond lengths differ significantly for the two
crystallographically independent anions at (3σ (Br(1)-F(1),
1.872(4) Å; Br(2)-F(2), 1.849(5) Å) and are attributable to
different Br-F‚‚‚N and Br-F‚‚‚O contacts between the NO+
cations and the fluorine ligands of the BrO3F2- anions (Figure
1c,d). Each fluorine atom of the crystallographically independent
BrO3F2- anions has four contacts with three NO+ cations such
that the Br-F‚‚‚[NO]3 coordination is a distorted tetrahedral
arrangement with different contacts to F(1) and F(2). Two cation
contacts to F(1) are of the N-O‚‚‚F type, with F‚‚‚O distances
of 2.746 and 2.854 Å, and N-O‚‚‚F angles of 132.1 and 143.8°,
(41) Pauling, L. The Nature of the Chemical Bond and the Structure of Molecules
and Crystals, 3rd ed.; Cornell University Press: Ithaca, NY, 1960; p 260.
(42) Bondi, A. J. Phys. Chem. 1964, 68, 441.
9
9418 J. AM. CHEM. SOC. VOL. 127, NO. 26, 2005

-
Synthesis of Salts Containing the BrO3F2 Anion
A R T I C L E S
Figure 1. (a) Structural unit and (b) packing diagram of [NO]2[BrO3F2][F]. The coordination environments of F(1), F(2), and F(3) are shown in (c), (d), and
(e), respectively. Thermal ellipsoids are shown at the 50% probability level.
results in elongation of the Br-F(1) bond. The N(1)-O(5)
(1.038(9) Å) and N(2)-O(6) (1.066(8) Å) bond lengths are
equivalent to within (3σ and therefore cannot be used to
correlate the O-N‚‚‚F interactions with Br-F(1) bond elonga-
tion. The existence of two discreet NO+ cations and their bond
length differences are, however, confirmed by the observation
of two N-O stretching frequencies in the Raman spectrum of
the salt (see Vibrational Spectra of [M][BrO3F2]).
which, in the absence of fluoride ions in the crystal lattice, is
expected to result in stronger ON+‚‚‚F-BrO3F contacts and
destabilization of the salt with respect to [NO]2[BrO3F2][F] and
BrO3F formation (eq 3).
2[NO][BrO3F2] f [NO]2[BrO3F2][F] + BrO3F (3)
(b) [N(CH3)4][BrO3F2]. The crystal structure of [N(CH3)4]-
[BrO3F2] (Figure S1) is consistent with an ionic formulation,
exhibiting well-separated and weakly interacting cations and
The fluoride ion (F(3)) of [NO]2[BrO3F2][F] is tetrahedrally
coordinated to four NO+ cations by means of long N‚‚‚F
contacts (2.216-2.412 Å; Figure 1e), which are significantly
longer than the N-F bond length of NOF (1.52 Å),43,44 but
shorter than the N‚‚‚F contacts to the BrO3F2- anions (2.544-
2.631 Å), reflecting the stronger Lewis basicity of the fluoride
ion. Although none of the O-N‚‚‚F contacts display the ideal
90° angle that maximizes the π-acceptor interaction with the
cation, the cations are coordinated side-on to the fluoride ion
with O-N‚‚‚F angles (102.3-111.4°) and O‚‚‚F(3) contact
distances (2.517-2.945 Å) that lie within or close to the sum
of the van der Waals radii for oxygen and fluorine (vide supra).
Coordination of the NO+ cations to the fluoride ion presumably
serves to stabilize the [NO]2[BrO3F2][F] salt by reducing the
+
anions. While the N(CH3)4 cations are fully ordered, the
-
BrO3F2 anions display a 2-fold (i.e., 90°) disorder along the
Br-O(1) axis, giving rise to two equally populated anion
orientations (Figure S1a). The disorder reflects the solid-state
packing arrangement of the salt, which can be described as
staggered columns of cations and anions running parallel to the
c-axis (Figure S1b). Whereas the disorder does not allow the
extraction of meaningful bond lengths and bond angles involving
disordered oxygen and fluorine atoms, the Br-O(1) bond length
(1.600(3) Å) is reliable and is in excellent agreement with the
Br-O distances determined for [NO]2[BrO3F2][F] (vide supra).
Although two possible positions for the disordered oxygen and
fluorine atoms can be resolved, the Br-O(2) (1.44(1) Å) and
Br-F(1) (1.77(6) Å) bond lengths are both artificially con-
tracted. The O(1)-Br-O(2) and O(1)-Br-F(1) bond angles,
which should be 120 and 90°, are 118.7(5) and 98.5(3)°,
respectively, and are less affected by the disorder.
-
cation interactions with the fluorine ligands of the BrO3F2
anions. This may account for failure to form [NO][BrO3F2]
(43) Magnuson, D. W. J. Chem. Phys. 1951, 19, 1071.
(44) Stephenson, C. V.; Jones, E. A. J. Chem. Phys. 1952, 20, 135.
9
J. AM. CHEM. SOC. VOL. 127, NO. 26, 2005 9419

A R T I C L E S
Lehmann and Schrobilgen
30
(c) Bond Length Comparisons among Known Br(VII)
Species. The nondisordered Br-O bond lengths obtained
from the crystal structures of [N(CH3)4][BrO3F2] and [NO]2-
[BrO3F2][F] are very similar to the average bond length of
and ν2(A1′), respectively, by analogy with XeO3F2 and
monomeric OsO3F2.31
The high-temperature phase, â-[Cs][BrO3F2], is easily dis-
tinguished from the R-phase on the basis of frequency shifts,
splittings of the Raman-active modes, and additional weak bands
in the Raman spectrum, which are assigned to the formally
Raman-inactive ν3(A2′′) and ν4(A2′′) modes. The ν1(A1′) (807
cm-1) and ν5(E′) (903 cm-1) BrO3 stretches are at somewhat
higher frequencies in the â-phase than in the R-phase. The ν5(E′)
band is also significantly broadened with shoulders at 899 and
908 cm-1, which are presumed to arise from site-symmetry
lowering and vibrational coupling within the unit cell. The
splittings of the ν6(E′) (386, 392 cm-1) and ν8(E′′) (404, 409
cm-1) modes are also consistent with the latter assumption. The
reduced symmetry, inferred from the splitting of the degenerate
BrO3 stretches and bends, also allows the observation of the
ν3(A2′′) and ν4(A2′′) modes. The BrO3 out-of-plane bend
(ν3(A2′′)) was assigned at 449 cm-1, while the higher frequency
bands (485, 497 cm-1) were assigned to the antisymmetric BrF2
stretch (ν4(A2′′)). The latter modes were observed at 468 and
504 cm-1 in the infrared spectrum of â-[Cs][BrO3F2] along with
the ν1(A1′) (800 cm-1) and ν5(E′) (881, 900, 909 cm-1) modes
(Figure 2b).
The spectra of [K][BrO3F2] and [Rb][BrO3F2] were similar
to that of â-[Cs][BrO3F2] and were assigned accordingly. In
contrast with [Cs][BrO3F2], only single phases were observed
for [K][BrO3F2] and [Rb][BrO3F2], which may reflect the higher
temperatures required for their syntheses. The symmetric and
antisymmetric BrO3 stretching frequencies of the K+ and Rb+
salts are slightly higher in frequency than in â-[Cs][BrO3F2]
and exhibit larger splittings of the ν5(E′) mode, and the splittings
of the ν2(A1′), ν6(E′), and ν8(E′′) modes are somewhat more
complex in the K+ and Rb+ salts than in â-[Cs][BrO3F2].
Although the splittings of the degenerate E′ modes are attribut-
able to site-symmetry lowering of the anion in the solid state,
the splittings of the nondegenerate ν2(A1′) mode in the K+ (420,
425 cm-1) and Rb+ (417, 427 cm-1) salts indicate that
vibrational coupling within their respective unit cells must also
occur. The formally Raman-inactive A2′′ modes appear as weak
bands in the Raman spectra of [K][BrO3F2] and [Rb][BrO3F2].
The asymmetric BrF2 stretching mode, ν4(A2′′), is factor-group
split and has similar band contours in the spectra of [K][BrO3F2]
(506, 523 cm-1), [Rb][BrO3F2] (497, 512 cm-1), and â-
[Cs][BrO3F2] (485, 497 cm-1). The BrO3 bend, ν3(A2′′), which
is a single band in [K][BrO3F2] (457 cm-1) and [Rb][BrO3F2]
(455 cm-1) but is split in â-[Cs][BrO3F2] (433, 449), is less
sensitive to the nature of the counterion.
BrO4 (1.603(16) Å)45-47 despite the higher coordination
-
number of BrO3F2-. As anticipated, the Br-O and Br-F bonds
are longer and more polar than the Br-F (1.708(3)) and Br-O
(1.582(1) Å) bond of neutral BrO3F48 and have bond orders
that are less than those of BrO3F (see Computational Results).
Although the average bond lengths for AsF6- (1.70(2) Å),49
+
SeF6 (1.69(1) Å),50 and BrF6 (1.666(11) Å)51 do not show a
significant contraction with increasing positive charge, the F‚‚‚F
bond orders51 and ligand packings of the XF6+ (X ) Cl, Br, I)
+
cations51,52 indicate that the fluorine ligands of BrF6 are not
close-packed. Ligand-ligand repulsions also do not appear to
strongly influence the bond lengths of lower-coordinate Br(VII)
species (vide supra). Thus, the anticipated trend of decreasing
Br-F bond length with increasing net positive charge is
observed for BrO3F2- (1.861(16) Å), BrO3F (1.708(3) Å),48 and
+
BrF6 (1.666(11) Å).51
Vibrational Spectra of [M][BrO3F2] (M ) K, Rb, Cs,
N(CH3)4) and [NO]2[BrO3F2][F]. Of the 12 fundamental modes
of vibration (Γvib ) 2A1′ + 2A2′′ + 3E′ + E′′) predicted for
BrO3F2- (Figure S2), the ν1(A1′), ν2(A1′), ν5(E′), ν6(E′), ν7(E′),
and ν8(E′′) modes are ideally only Raman-active and the ν3(A2′′),
ν4(A2′′), ν5(E′), and ν6(E′) modes are ideally only infrared-active.
The Raman spectra (-163 °C) of R-[Cs][BrO3F2], â-[Cs]-
[BrO3F2], [Rb][BrO3F2], [K][BrO3F2], [N(CH3)4][BrO3F2], and
[NO]2[BrO3F2][F] and the room-temperature infrared spectrum
of â-[Cs][BrO3F2] are shown in Figure 2. The vibrational
-
assignments for the BrO3F2 anion (Table 3) were made with
the assistance of quantum mechanical calculations in Tables 3
and S1 (also see Computational Results and Supporting
Information) and by comparison with the vibrational spectra of
XeO3F2 and OsO3F2.31,32 Trends in the BrO3 and BrF2
30
stretching frequencies reflect the anticipated cation-anion
interactions in the salt, paralleling the cation charge-to-radius
ratios,27 which increase in the order N(CH3)4 < Cs+ < Rb+
+
< K+ (see Supporting Information).
(a) [M][BrO3F2] (M ) Cs, Rb, K). The low-temperature
R-phase of [Cs][BrO3F2] exhibits six Raman-active bands,
consistent with D3h symmetry. The most intense band in the
spectrum (802 cm-1) was assigned to the symmetric BrO3
stretch, ν1(A1′), and the band at 896 cm-1 to the asymmetric
BrO3 stretch, ν5(E′). These assignments are supported by the
calculated frequencies and the anticipated trend, νs < νas,
although it is noteworthy that the order of the ν1(A1′) and ν5(E′)
(b) [N(CH3)4][BrO3F2]. The Raman spectrum of [N(CH3)4]-
31,32
[BrO3F2] is characteristic for a BrO3F2- anion and a N(CH3)4
+
frequencies is reversed for OsO3F2
(see Symmetry Force
cation.38,53-55 The BrO3F2 anion disorder (see X-ray Crystal
-
Constants Derived from B-Matrix Analyses). The bands at 229,
383, 414, and 433 cm-1 were assigned to ν7(E′), ν6(E′), ν8(E′′),
Structures of [NO]2[BrO3F2][F] and [N(CH3)4][BrO3F2]) pre-
vents an unambiguous correlation between the free ion (D3h),
site (C4V), and crystallographic unit cell (D4h) symmetries;
however, the observed broadening or splitting of the ν2(A1′),
ν5(E′), ν6(E′), ν7(E′), and ν8(E′′) modes is indicative of factor-
group splitting. The BrO3 stretches occur at 801 cm-1 (ν1(A1′))
and 896 cm-1 (ν5(E′)), with the former again being the most
(45) Siegel, S.; Tani, B.; Appelman, E. Inorg. Chem. 1969, 8, 1190.
(46) Gebert, E.; Peterson, S. W.; Reis, A. H.; Appelman, E. H. J. Inorg. Nucl.
Chem. 1981, 43, 3085.
(47) Gallucci, J. C.; Gerkin, R. E.; Reppart, W. J. Acta Crystallogr. 1989, C45,
701.
(48) Appelman, E. H.; Beagley, B.; Cruickshank, D. W. J.; Foord, A.; Rustad,
S.; Ulbrecht, V. J. Mol. Struct. 1976, 35, 139.
(49) Lehmann, J. F.; Dixon, D. A.; Schrobilgen, G. J. Inorg. Chem. 2001, 40,
3002.
(50) Ewing, V. C.; Sutton, L. E. Trans. Faraday Soc. 1963, 59, 1241.
(51) Lehmann, J. F.; Schrobilgen, G. J.; Christe, K. O.; Kornath, A.; Suontamo,
R. J. Inorg. Chem. 2004, 43, 6905.
(53) Berg, R. W. Spectrochim. Acta 1978, 34A, 655.
(54) Kabisch, G.; Klose, M. J. Raman Spectrosc. 1978, 7, 311.
(55) Mercier, H. P. A.; Sanders, J. C. P.; Schrobilgen, G. J. J. Am. Chem. Soc.
1994, 116, 2921.
(52) Robinson, E. A.; Gillespie, R. J. Inorg. Chem. 2003, 42, 3865.
9
9420 J. AM. CHEM. SOC. VOL. 127, NO. 26, 2005

-
Synthesis of Salts Containing the BrO3F2 Anion
A R T I C L E S
Figure 2. Raman spectra (recorded at -163 °C) of (a) R-[Cs][BrO3F2]; (b) â-[Cs][BrO3F2], the band at 802 cm-1 (‡) arises from residual R-[Cs][BrO3F2]
and the upper trace is the room-temperature infrared spectrum of â-[Cs][BrO3F2] in an AgCl pellet; (c) [Rb][BrO3F2]; (d) [K][BrO3F2]; (e) [N(CH3)4]-
+
[BrO3F2], unlabeled bands arise from the N(CH3)4 cation and the band at 460 cm-1, tentatively assigned to ν4(A2′′) of BrO3F2-, may alternatively be
assigned to, or overlap with, ν19(F2) of N(CH3)4+; (f) [NO]2[BrO3F2][F]. Asterisks (*) denote FEP sample tube lines, and daggers (†) denote instrumental
artifacts.
intense feature in the spectrum and the latter exhibiting some
broadening. The frequencies of the remaining modes and their
relative intensities are also in good agreement with those of
the alkali metal salts. The formally Raman-inactive ν3(A2′′) was
not observed, and the band at 460 cm-1 is tentatively assigned
to the formally Raman-inactive ν4(A2′′) mode of BrO3F2-, which
which are expected to be significantly lower in frequency than
those of ONF (492 and 735 cm-1, respectively)56 as a result of
the long N‚‚‚F contact distances and are broadened because there
are several such contacts (Figure 1 and Table 2). The presence
of two bands in the N-O stretching region (2203, 2246 cm-1
)
is in accord with the two crystallographically unique cation sites
+
may overlap with ν19(F2) of the N(CH3)4 cation.38,53-55
(Figure 1c,d). Although the N-O stretching frequencies agree
(c) [NO]2[BrO3F2][F]. The Raman spectrum of [NO]2-
more closely with that of NO+ (2273 cm-1 57
) than with that of
[BrO3F2][F] is similar to that of the alkali metal and N(CH3)4
gaseous ONF (1844 cm-1),58-60 the low-frequency shift is
+
salts, with the added exception that the low-frequency region
(100-340 cm-1) is dominated by broad bands of medium
intensity. It is likely that a number of these bands arise from
O-N‚‚‚F(1,2,3) bending and ON‚‚‚F(1,2,3) stretching modes,
(56) Smardzewski, R. R.; Fox, W. B. J. Am. Chem. Soc. 1974, 96, 304.
(57) Babaeva, V. P.; Rosolovskii, V. Y. Russ. J. Inorg. Chem. 1971, 16, 471.
(58) Jones, E. A.; Woltz, P. J. H. J. Chem. Phys. 1950, 18, 1516.
(59) Magnuson, D. W. J. Chem. Phys. 1952, 20, 380.
(60) Jones, L. H.; Asprey, L. B.; Ryan, R. R. J. Chem. Phys. 1967, 47, 3371.
9
J. AM. CHEM. SOC. VOL. 127, NO. 26, 2005 9421

A R T I C L E S
Table 3. Vibrational Frequencies, Intensities, and Assignments for the XO3F2 (X ) Cl, Br) Anions
Lehmann and Schrobilgen
-
-a
-
BrO3F2
ClO3F2
Raman
IR
Raman
c
d
e
+ f
assignment (D3h
)
calcdb
K+
Rb+
R
-Cs+
â
-Cs+
â
-Cs+
[NO+]2[F-
]
N(CH3)4
calcdb
ν1(A1′)
ν2(A1′)
νs(XO3)
νs(XF2)
788(100)[0]
435(45)[0]
811(100)
435(6)
420(13)
523(2)
506(3)
457(5)
810(100)
427(5)
417(22)
512(3)
497(3)
455(5)
802(100)
433(19)
807(100)
421(sh)
417(17)
497(3)
485(3)
449(6)
433(3)
908(sh)
903(7)
899(sh)
392(14)
386(7)
383(sh)
227(12)
800[w,br]
806(100)
436(21)
425(28)
801(100)
442(14)
434(sh)
875(100)[0]
364(40)[0]
ν3(A2′′)
ν4(A2′′)
ν5(E′)
νas(XF2)
529(0)[389]
447(0)[17]
504[m]
468[m]
630(0)[376]
459(0)[317]
1144(24)[614]
δ(XO3)
oop
νas(XO3)
460(6)g
896(5)
883(32)[304]
914(5)
908(5)
904(7)
397(19)
393(sh)
912(3)
907(sh)
904(8)
394(20)
386(7)
381(3)
228(sh)
231(11)
403(11)
896(8)
909[s]
900[s]
881[sh]
917(3)
908(7)
900(6)
397(10)
390(5)
381(7)
240(sh)
233(sh)
415(14)
407(25)
ν6(E′)
δ(XO3)
ip
378(26)[76]
383(23)
395(3)
389(4)
382(2)
228(1)
222(1)
414(12)
408(11)
513(29)[54]
ν7(E′)
δ(XF2)
Fr(XF2)
207(6)[<1]
227(6)
233(6)
406(11)
229(13)
414(27)
247(5)[4]
ν8(E′′)
384(24)[0]
409(18)
404(15)
456(19)[0]
a All experimental frequencies are from Raman spectra, except for â-[Cs][BrO3F2] for which infrared frequencies are also provided. Frequencies are
given in cm-1. All Raman spectra were recorded at -163 °C in FEP sample tubes. The experimental intensities, given in parentheses, are scaled relative to
the intensity of the ν1 mode, which is assigned a value of 100. The abbreviation, sh, denotes a shoulder. b MPW1PW91/DZVP. The calculated absolute
-
Raman intensities in amu Å-4 may be obtained by dividing the relative calculated intensities by the following scaling coefficients: 2.45 for BrO3F2 and
-
2.38 for ClO3F2
.
c The symbols ν, δ, and Fr denote stretch, bend, and rock, respectively. The abbreviations s, as, oop, and ip denote symmetric, antisymmetric,
out-of-plane, and in-plane, respectively. d The infrared spectrum of â-[Cs][BrO3F2] was obtained in a AgCl pellet at ambient temperature. The abbreviations
w, m, s, and br denote weak, medium, strong, and broad infrared lines, respectively. e The NO+ stretches occur at 2203(8) and 2246(31) cm-1. Additional
broad bands were observed at 340(6), 302(6), 295(8), 271(11), 255(10), 233(11), 217(13), 205 sh, 121(23), 100(68), 65(7) cm-1 and are tentatively assigned
to O-N‚‚‚F(1,2,3) bending and ON‚‚‚F(1,2,3) stretching modes associated with NO+ and BrO3F2 interactions in the crystal lattice and to lattice modes.
-
f Bands at 3042(16), 2995(6), 2919(5), 2881(3), 2866(2), 2809(5), 1469(11), 1466(13), 1412(5), 1286(3), 1184(2), 1177(3), 949(19), 755(14), 460(6), and
329(2) cm-1 were assigned to N(CH3)4+ by comparison with previous assignments given in refs 38, 54-56. g This band may result from the overlap of the
+
formally Raman-inactive ν4(A2′′) mode and ν19(F2) of the N(CH3)4 cation.
consistent with significant interactions between the NO+ cations
the HF, MP2, and LDF (MPW1PW91) (Table S3) methods are
comparable, only the LDF results are referred to in the present
discussion. The calculated geometries of the parent oxide
fluorides, BrO3F and ClO3F, were used as benchmarks for
comparison of the effects of fluoride ion coordination on the
X-O and X-F (X ) Cl, Br) bond lengths. The energy-
and the BrO3F2 and F- anions. The splittings of the ν2(A1′),
-
ν5(E′), ν6(E′) ν7(E′), and ν8(E′′) modes of BrO3F2- in the Raman
spectrum of [NO]2[BrO3F2][F] are consistent with factor-group
splitting (see complementary discussion in Supporting Informa-
tion and Table S2).
1
NMR Spectroscopy. The 19F, H, and 13C NMR spectra of
2-
minimized gas-phase structures of fac-BrO3F3 and mer-
2-
[N(CH3)4][BrO3F2] dissolved in CH3CN at -40 °C were
obtained from a sample that had been stored at -196 °C and
had not been previously warmed above -40 °C. The 19F NMR
spectrum exhibited a single resonance at 237.0 ppm (∆ν1/2, 360
Hz), which is assigned to the chemically equivalent fluorines
BrO3F3 are also provided in Table S3 and are discussed in
the Supporting Information.
The geometry of the BrO3F2- anion in [NO]2[BrO3F2][F] (see
X-ray Crystal Structures of [NO]2[BrO3F2][F] and [N(CH3)4]-
[BrO3F2]) exhibited only small distortions from the ideal trigonal
bipyramidal geometry, justifying the calculation of BrO3F2- and
of the BrO3F2 anion. The 19F chemical shift of BrO3F2 is
-
-
+
significantly more shielded than that of BrF6 (339,3 3373,61
-
ClO3F2 constrained to D3h symmetry. The calculated Br-O
ppm) and BrO3F (269-274 ppm),24 with the general trend of
increased deshielding with increasing formal charge being
adhered to for these Br(VII) species. The broad line width is
attributed to the nearly completely quadrupole-collapsed
1J(79,81Br-19F) spin-spin coupling, as is the case for BrO3F
(∆ν1/2, 70-217 Hz).24 Quadrupolar collapse in these lower
symmetry species contrasts with that of the highly symmetric
bond length (1.639 Å) is moderately longer than the average
experimental value (1.603(6) Å); however, the difference
-
between the Br-O bond lengths of BrO3F2 and BrO3F is
comparable when determined by theory (0.03 Å) and by
experiment (0.01-0.03 Å). The calculated Br-F bond length
(1.905 Å) is also slightly longer than the average obtained from
the structure of [NO]2[BrO3F][F]; however, in this instance the
difference between the values calculated for BrO3F2- and BrO3F
(0.10 Å) is less than that observed experimentally (0.15 Å).
The difference is attributed to significant ON‚‚‚F-BrO3F
interactions in the solid state that further elongate the Br-F
bonds of the anion.
+
BrF6 cation, for which the near-zero electric field gradient at
the octahedrally coordinated quadrupolar 79Br and 81Br (I )
3/2) nuclei results in slow quadrupolar relaxation and observation
of the 1J(79,81Br-19F) couplings.3,51 The 1H and 13C resonances
of the N(CH3)4 cation were observed at 3.037 and 55.4 ppm,
+
1
respectively, with a J(13C-1H) ) 143 Hz.
Although attempts to prepare [N(CH3)4][ClO3F2] were unsuc-
cessful, the anion is predicted to be thermodynamically stable
with respect to dissociation to ClO3F and F- in the gas phase
(vide infra). The weak fluoride ion acceptor properties of ClO3F
are reflected in the Cl-F bond lengths of the anion (1.881 Å),
Computational Results. (a) BrO3F2- and ClO3F2- Geom-
etries. Although the energy-minimized gas-phase structures of
-
BrO3F, ClO3F, BrO3F2-, and ClO3F2 determined by use of
(61) Schroer, T.; Christe, K. O. Inorg. Chem. 2001, 40, 2415.
9
9422 J. AM. CHEM. SOC. VOL. 127, NO. 26, 2005

-
Synthesis of Salts Containing the BrO3F2 Anion
A R T I C L E S
which were calculated to be ca. 0.20 Å longer than that for
ClO3F (1.685 Å), in accord with the smaller difference (ca. 0.10
Å) obtained for the bromine analogues. Additional elongation
synthesize [N(CH3)4]2[BrO3F3] (see Attempted Syntheses of
[N(CH3)4]2[BrO3F3] and [N(CH3)4][ClO3F2] in the Experimental
Section).
-
(c) Vibrational Frequencies of BrO3F2- and ClO3F2-. The
and destabilization of the Cl-F bonds of ClO3F2 relative to
-
vibrational frequencies and intensities of BrO3F2- and ClO3F2
that of ClO3F would be expected in the solid state, where
contacts between the fluorine ligands and countercation are
likely and may account, in part, for the inability to prepare salts
of this anion. The difference (ca. 0.03 Å) between the calculated
were calculated by use of HF, MP2, and LDF (MPW1PW91)
(Table S1) methods and were used to assist in the assignment
-
of the experimental vibrational spectra of the BrO3F2 salts.
-
The LDF method provided the most reliable vibrational
frequencies and intensities for the ClO3F2- and BrO3F2- anions,
which are referred to in the present discussion and in Table 3.
Cl-O bond lengths of ClO3F2 (1.465 Å) and ClO3F (1.438
-
Å) is similar to that calculated for BrO3F2 and BrO3F.
(b) Enthalpies of Fluoride Ion Attachment to BrO3F and
ClO3F. The enthalpies of fluoride ion attachment for BrO3F,
BrO3F2-, and ClO3F have been determined by the G2 method
(Table S4), which includes fourth order (MP4) and configuration
interaction (QCISD) energy corrections. The method differs
from the near-local density functional (NLDF)26 and MP2/PDZ62
methods used previously. The accuracy of the G2 method was
verified by calculating the enthalpies of fluoride ion attachment
for HF, COF2, SO2, BF3, and AlF3 (Table S4), which have
accurately determined experimental values and which have been
used as benchmarks in previous computational studies.26,62
(d) Force Constants for BrO3F2-, ClO3F2-, XeO3F2, and
OsO3F2. Internal force constants have been previously deter-
mined for XeO3F230,70-72 and OsO3F271-73 by use of the Wilson
-
FG-matrix method74 and are now calculated for the BrO3F2
anion by the same method. To complement these studies and
to allow for comparisons with the unknown ClO3F2- anion, the
symmetry force constants of BrO3F2-, ClO3F2-, XeO3F2, and
OsO3F2 have been calculated by the B-matrix method,75 which
provides a fully determined force field that includes off-diagonal
components.
(i) Internal Force Constants Derived from General Va-
lence Force Field (GVFF) Analyses. The internal force
constants and potential energy distributions of BrO3F2-, XeO3F2,
and OsO3F2, determined by the FG-matrix method, are sum-
marized in Tables 4 and S5, respectively. The lower fD value
Both BrO3F and ClO3F are predicted to be fluoride ion
acceptors in the gas phase; however, the enthalpy of fluoride
ion attachment is significantly more exothermic for BrO3F
(-261.1 kJ mol-1) than for ClO3F (-132.3 kJ mol-1). The
enthalpy of fluoride ion attachment to BrO3F, determined in
the present study, is 11% less than the value reported earlier
-
of BrO3F2 with respect to that of XeO3F2 is attributed to the
higher formal oxidation state of xenon and resulting lower
polarities of the Xe-O bonds. Despite the neutral charges and
formal +8 oxidation states of osmium and xenon, fD is
significantly greater for OsO3F2 than for XeO3F2. The difference
likely reflects the relative degrees of dπ-pπ bonding between
the central atom and the oxygen ligands, which is expected to
be significant for d0 species, but a minor contributor in the
bonding of the main-group species.76 The greater Os-O bond
strength of OsO3F2 is also indicated by the calculated Os-O
bond length (1.717 Å), which is ca. 0.08 Å shorter than the
calculated Xe-O bond length of XeO3F2 (1.802 Å) even though
xenon is in row five and osmium is in row six of the periodic
table.
(-292 kJ mol-1 26
but is bracketed by the values previously
)
calculated for XeOF4 (-267 kJ mol-1), TeF6 (-257 kJ mol-1),
and XeF4 (-239 kJ mol-1).26 In view of the experimentally
established stabilities of XeOF5 , , ,
- 63 TeF7- 64 and XeF5- 65 it is
-
not surprising that salts of the BrO3F2 anion have proven to
be isolable.
The enthalpy of fluoride ion attachment for ClO3F determined
by the G2 method is also lower than that obtained by the NLDF
method (-180 kJ mol-1).26 The present study indicates that the
fluoride ion acceptor strength of ClO3F is intermediate with
respect to those of PF3 (-187.9 kJ mol-1), HF (-154.0 kJ
mol-1), NO2F (-80.3 kJ mol-1), and NOF (-72.8 kJ mol-1
)
in the gas phase.62 Although salts of PF4
and HF2
and NOF2 are unknown. The
are
-66,67
-68
The Xe-F and Os-F stretching force constants (fd) obtained
for XeO3F2 and OsO3F2 in the present work are in good
-69
-
known, those of NO2F2
-
ranking of ClO3F2 within this series is consistent with the
agreement with the previously reported values determined by
71,72
and OsO3F2,71 but are
minimum enthalpy of fluoride ion attachment (-126 to -146
the FG-matrix method for XeO3F2
kJ mol-1 26
) generally required for the formation of a stable salt,
somewhat greater than those reported in refs 72 and 73 for
OsO3F2. The trends among the fd values for BrO3F2-, XeO3F2,
and OsO3F2 are similar to those of fD and are also attributable
to net charge, formal oxidation state of the central atom, and
relative degrees of dπ-pπ bonding. The smaller difference
and inability to prepare [N(CH3)4][ClO3F2]. The fluoride ion
-
acceptor properties of BrO3F2 were also investigated by the
G2 method. The positive enthalpies of formation arrived at for
the gas-phase formation of fac-BrO3F32- (362.1 kJ mol-1) and
2-
mer-BrO3F3 (367.8 kJ mol-1) indicate that both isomers are
between the fd values of XeO3F2 and OsO3F2 (0.83 mdyne Å-1
)
when compared with that of their fD values (1.37 mdyne Å-1
)
unstable and is consistent with the present failed attempts to
is consistent with the reluctance of fluorine to form multiple
bonds.
(62) Christe, K. O.; Dixon, D. A.; McLemore, D.; Wilson, W. W.; Sheehy, J.
A.; Boatz, J. A. J. Fluorine Chem. 2000, 101, 151.
(63) Christe, K. O.; Dixon, D. A.; Sanders, J. C. P.; Schrobilgen, G. J.; Tsai, S.
S.; Wilson, W. W. Inorg. Chem. 1995, 34, 1868.
(70) Natarajan, A.; Chockalingam, K. Pramana 1977, 9, 573.
(71) Gnanasekaran, S.; Ranganayaki, S.; Gnanasekaran, P.; Sampath Krishnan,
S. Asian J. Chem. 1989, 1, 173.
(72) Mohan, S.; Vasuki, G. Proc. Natl. Acad. Sci., India, Sect. A 1989, 59, 163.
(73) So, S. P. Z. Phys. Chem., Neue Folge 1978, 109, 157.
(74) Wilson, E. B. J. Chem. Phys. 1941, 9, 76.
(75) Komornicki, A.; BMATRIX, version 2.0; Polyatomics Research Institute:
Palo Alto, CA, 1996.
(76) Reed, A. E.; Schleyer, P. v. R. J. Am. Chem. Soc. 1990, 112, 1434.
(64) Selig, H.; Sarig, S.; Abramowitz, S. Inorg. Chem. 1974, 13, 1508.
(65) Christe, K. O.; Curtis, E. C.; Dixon, D. A.; Mercier, H. P. A.; Sanders, J.
C. P.; Schrobilgen, G. J. J. Am. Chem. Soc. 1991, 113, 3351.
(66) Wermer, P.; Ault, B. S. Inorg. Chem. 1981, 20, 970.
(67) Christe, K. O.; Dixon, D. A.; Mercier, H. P. A.; Sanders, J. C. P.;
Schrobilgen, G. J.; Wilson, W. W. J. Am. Chem. Soc. 1994, 116, 2850.
(68) Gilbert, A. S.; Sheppard, N. Spectrochim. Acta, Part A 1976, 32, 923.
(69) Lawlor, L.; Passmore, J. Inorg. Chem. 1979, 18, 2923.
9
J. AM. CHEM. SOC. VOL. 127, NO. 26, 2005 9423

A R T I C L E S
Lehmann and Schrobilgen
-
Table 4. Internal and Symmetry Force Constants for BrO3F2-, XeO3F2, OsO3F2, and ClO3F2 and Potential Energy Distributions (PED)
Derived from B-Matrix Analyses
internal force constants (GVFF)a,b
symmetry force constants (B-matrix)b,c
-
-
-
BrO3F2
XeO3F2
OsO3F2
ClO3F2
BrO3F2
XeO3F2
OsO3F2
fD
fd
fâ
fR
5.86
2.15
1.63
0.59
0.11
0.28
0.003
0.04
0.04
0.12
0.12
-0.08
0.11
-0.02
-0.02
-0.10
0.13
6.27
3.29
1.36
0.66
-0.07
0.38
-0.0003
0.005
0.005
-0.13
0.20
-0.19
0.03
0.01
0.01
7.64
4.12
0.94
0.46
0.19
0.49
0.006
0.16
0.06
0.09
-0.12
-0.37
0.01
-0.01
-0.03
-0.02
-0.06
0.005
A1′ species:
F11 ) fD + 2fDD
6.98
1.68
0.43
5.83
2.09
0.28
5.45
2.92
-0.09
8.73
4.30
0.49
F22 ) fd + 2fdd
F12 ) 61/2fDd
fDD
fdd
fDd
fDâ
fDR
fdâ
fââ
fRR
fRâ
f D′ â
f D′ R
f d′â
A2′′ species:
F33 ) fd - fdd
1.73
2.03
0.78
2.19
1.65
0.41
3.04
1.28
0.11
4.13
0.85
0.61
F44 ) fâ - fââ - 2f â′ â - 2f ′′
ââ
F34 ) (3(fdâ - f ′ ))1/2
dâ
E′ species:
F55 ) fD - fDD
6.97
0.85
5.64
0.59
5.64
0.43
7.63
0.64
F66 ) fR - fRR
F77 ) fâ + fââ - f â′ â - f ′′
1.57
1.36
1.10
1.18
ââ
F56 ) f ′ - fDR
-0.08
0.24
0.02
0.19
0.06
0.06
0.14
0.15
DR
F57 ) 21/2(fDâ - f ′ )
F67 ) -21/2(fRâ -Dfâ′ )
-0.37
-0.32
-0.26
-0.28
Râ
0.15
-0.22
-0.01
E′′ species:
f ′
F88 ) fâ - fââ - f â′ â - f ′′
1.40
1.14
0.85
1.14
ââ
ââ
f ′
-0.09
Râ
-
-
ClO3F2
BrO3F2
PED (B-matrix)g
XeO3F2
PED (B-matrix)g
OsO3F2
PED (B-matrix)g
assgnt
DFT freqd,f
PED (B-matrix)g
DFT freqd,f
DFT freqd,f
DFT freqd,f
ν1(A1′)
ν2(A1′)
ν3(A2′′)
ν4(A2′′)
ν5(E′)
862
383
622
474
1133
100 S1
100 S2
77 S3 + 23 S4
1 S3 + 99 S4
12 S5/6 + 57 S6/5
22 S7/8+ 8 S9/10
787
430
529
423
886
100 S1
100 S2
80 S3 + 20 S4
4 S3 + 96 S4
76 S5/6 + 13 S6/5
761
510
595
326
845
100 S1
100 S2
93 S3 + 7 S4
1 S3 + 99 S4
93 S5/6 + 3 S6/5
3 S8/7 + 1 S10/9
965
617
655
262
957
99 S1 + 1 S2
99 S2 + 1 S1
100 S3
2 S3 + 98 S4
21 S5/6 + 76 S6/5
1 S7/8 + 2 S8/7
+
+
+
+
5 S7/8 + 3 S8/7
+
2 S9/10 + 1 S10/9
3 S7/8 + 77 S8/7
20 S10/9
ν6(E′)
484
1 S5/6 + 85 S 8/7
14 S10/9
+
357
+
276
74 S8/7 + 26 S10/9
312
85 S8/7 + 15 S10/9
ν7(E′)
ν8(E′′)
235
435
51 S8/7 + 49 S10/9
100 S11/12
193
362
64 S8/7 + 36 S10/9
100 S11/12
157
291
71 S8/7 + 29 S10/9
100 S11/12
189
353
51 S8/7 + 49 S10/9
99 S11/12 + 1 S12/11
a The internal force constants corresponding to E-O and E-F stretching are denoted by fD and fd, respectively, and the O-E-O and O-E-F bends are
denoted by fR and fâ, respectively. The interaction force constants are denoted by fxx, fxy, f ′ , and f ′ , where x and y are D, d, R, or â. The relative
xy
contributions of the internal coordinates to the PED of a normal coordinate are multiplied by 1x0x0. b Force constants for bond stretches and bends are given
in mdyne Å-1 and mdyne Å rad-2, respectively. c The force constants, Fii (i ) 1‚‚‚8), correspond to the normal modes, νi. The symmetry force constants Fij
correspond to the off-diagonal interactions between νi and νj, where i * j. d Frequencies are given in cm-1 e Frequencies calculated from the normal coordinate
.
analysis. f The vibrational frequencies and second derivative matrix used for the determination of the symmetry force constants were calculated by use of the
-
SVWN method in conjunction with DZVP basis sets. When constrained to D3h symmetry, the energy-minimized E-F and E-O bond lengths were BrO3F2
(1.916, 1.646 Å), XeO3F2 (1.983, 1.802 Å), OsO3F2 (1.897, 1.717 Å), and ClO3F2- (1.878, 1.477 Å), respectively. g The PEDs are given in percentages and
have been normalized. The following symmetry coordinates were used and appear in the B-matrix potential energy distributions: S1 ) E-O1 + E-O2 +
E-O3, S2 ) E-F1 + E-F2, S3 ) E-F1 - E-F2, S4 ) O1EF1 + O2EF1 + O3EF1 - O1EF2 - O2EF2 - O3EF2, S5 ) 2E-O1 - (E-O2
+ E-O3), S6 ) E-O2 - E-O3, S7 ) 2 O1EO3 - ( O3EO2 + O1EO2), S8 ) O3EO2 - O1EO2, S9 ) 2( O2EF1 + O2EF2) - ( O1EF1 +
O3EF1 + O1EF2 + O3EF2), S10
+
)
O1EF1 + O1EF2 - O3EF2 - O3EF2, S11 ) 2( O2EF1 - O2EF2) - O1EF1 + O1EF2 - O3EF1
O3EF2, S12
)
O1EF1 - O1EF2 - O3EF1 + O3EF2. Their explicit forms are given in refs 72 and 77.
-
The O-E-F bending force constants, fâ, of BrO3F2
,
are listed in Table 4. The relationships between the symmetry
coordinates and internal coordinates of these systems are given
in refs 72 and 77, and descriptions of the symmetry coordinates
are provided in Table 4 (footnote g). The relationships between
the symmetry force constants and the internal force constants
of trigonal bipyramidal EY3Z2-type molecules have been derived
previously72 and are also provided in Table 4. A full comparison
of the force constants obtained by the two methods is not
possible because the FG-matrix method does not properly
address all of the off-diagonal terms, which are arbitrarily set
equal to zero. The values of diagonal terms (F11/F55 and fD;
F22/F33 and fd, F66 and fR; F77/F88 and fâ) for BrO3F2-, XeO3F2,
and OsO3F2 are in general agreement, which is consistent with
the relatively small values of the internal interaction force
constants in the expressions for symmetry force constants.
The potential energy distributions determined from the
B-matrix analyses of ClO3F2-, BrO3F2-, XeO3F2, and OsO3F2
(Table 4) reveal that there is no mixing of the symmetry
coordinates for the A1′ modes. Among the A2′′ modes,
XeO3F2, and OsO3F2 are large relative to the O-E-O bending
force constants, fR. The repulsive interligand O‚‚‚F interactions
in these species may account, in part, for the greater magnitudes
of these force constants, but the observation that fâ is greater
for XeO3F2 than for OsO3F2, despite the shorter bond lengths
and ligand-ligand distances calculated for OsO3F2, implies that
other factors are involved. The decrease in fâ with increasing
principal quantum number of the central element over the series
suggests that the trend may be related to valence orbital
diffuseness. A similar, but less pronounced, trend occurs among
the O-E-O bending force constants of BrO3F2-, XeO3F2, and
OsO3F2, whose overall lower values are attributed to the large
O-E-O bond angles (120°) that result in decreased interligand
O‚‚‚O interactions.
(ii) Symmetry Force Constants Derived from B-Matrix
Analyses. The unscaled symmetry force constants and potential
energy distributions of BrO3F2-, XeO3F2, OsO3F2, and the
-
unknown ClO3F2 anion determined by the B-matrix method
9
9424 J. AM. CHEM. SOC. VOL. 127, NO. 26, 2005

-
Synthesis of Salts Containing the BrO3F2 Anion
A R T I C L E S
significant mixing only occurs between the asymmetric EF2
stretches, S3, and the out-of-plane umbrella bends, S4, of ν3(A2′′)
in ClO3F2- and BrO3F2-. The three E′ modes have contributions
from S5-S10, with the degree of coupling being least for ν5(E′).
The ν5(E′) mode is best described in terms of the degenerate
EO3 stretching symmetry coordinates, S5/6, for BrO3F2-, XeO3F2,
and OsO3F2. In the case of ClO3F2-, they are extensively mixed
with the EO3 bending symmetry coordinates, S7/8. In contrast,
ν6(E′) and ν7(E′) exhibit extensive mixing among their EOF
bending coordinates, S7/8 and S9/10, but little or no mixing with
the S5/6 stretching coordinates.
The positive charges on the bromine (2.71-2.78) and chlorine
(2.43-2.49) atoms exhibit little variation among the oxygen-
-
rich species XO4-, XO3F, and XO3F2 (X ) Br, Cl) and are
similar to those reported for BrF6+ (2.86) and ClF6+ (2.51) using
the same computational method.51 The positive charges on the
chlorine atoms are approximately 0.3 less than those on the
bromine atoms of their bromine analogues and are attributed to
the higher electronegativity of chlorine. The relative invariance
of the central halogen atom charge requires that the atomic
charges on the oxygen and fluorine ligands compensate the net
charges of these species. The negative charge distributions
among the oxygen and fluorine ligands of XO3F (82/82% O,
Although the E-O bonds of XeO3F2 are expected to be less
polar than those of BrO3F2-, their symmetric (F11) and anti-
symmetric (F55) EO3 stretching force constants are similar. This
may, in part, arise from the greater underestimation of the EO3
stretching frequencies (average νexpt - νcalcd values for ν1(A1′)
-
18/18% F) and XO3F2 (63/67% O, 38/33% F) are similar for
X ) Cl and Br, with the negative charge distribution shifting
toward the fluorine atoms in the anions, in accord with the
enhanced polarities and lower bond orders of the X-F bonds
(vide infra). A comparison of the calculated X-F bond orders
of XO3F and XO3F2- shows that they correlate with the relative
fluoride ion affinities of BrO3F and ClO3F, with a Br-F bond
order for BrO3F2- that is 73% of that calculated for BrO3F and
and ν4(A2′′) are given in parentheses) of XeO3F2 (49 cm-1
)
-
relative to those of BrO3F2 (12 cm-1). The main-group F11
and F55 values contrast with significantly higher values for
OsO3F2, which may be attributed to stronger Os-O bonds
arising from dπ-pπ bonding contributions, as already noted
for the internal force constants. Comparisons of the EO3
-
a Cl-F bond order for ClO3F2 that is only 58% of that
calculated for ClO3F. Although the calculated Cl-O bond orders
are slightly higher (ca. 5%), the X-O bond orders show the
same ordering and little variation for each pair of analogues,
XO4-, XO3F, and XO3F2-. Similar trends are found among the
ligand valencies of these species (Table S6).
-
-
stretching force constants of BrO3F2 with those of ClO3F2
reveal that F11 and F55 are 1.15 and 1.33 mdyne Å-1 higher,
respectively, for ClO3F2-, in accord with the calculated Cl-O
bond order of ClO3F2-, which is 7% greater than the Br-O
bond order of BrO3F2- (Table S6) and the weaker fluoride ion
acceptor properties of ClO3F. The value of F11 is similar to that
of F55 for ClO3F2-, BrO3F2-, and XeO3F2, but it is significantly
greater than F55 for OsO3F2, accounting for the occurrence of
the symmetric OsO3 stretch, ν1(A1′), at higher frequency than
the antisymmetric OsO3 stretch, ν5(E′).31
Conclusion
Perbromyl fluoride is a sufficiently strong fluoride ion
acceptor toward MF (M ) K, Rb, Cs, N(CH3)4), and NOF to
-
form stable salts of the BrO3F2 anion. With the exception of
[NO]2[BrO3F2][F], which has a significant dissociation vapor
pressure, the BrO3F2- salts studied are all stable under dynamic
vacuum at -40 °C, and â-[Cs][BrO3F2] can be handled for up
to several hours under inert conditions at ambient temperatures.
The fully assigned vibrational spectra of the aforementioned
salts and X-ray crystal structures of [NO]2[BrO3F2][F] and
[N(CH3)4][BrO3F2] establish the trigonal bipyramidal geometry
The internal force constant, fd, dominates the expressions for
the symmetric (F22) and antisymmetric (F33) EF2 symmetry force
constants, accounting for similar trends among the internal and
symmetry force constants associated with the symmetric and
antisymmetric EF2 stretches. The symmetric EF2 stretching
symmetry force constant is greater than its antisymmetric
counterpart for OsO3F2 (4.30, 4.13 mdyne Å-1), XeO3F2 (2.92,
-
of BrO3F2 predicted by the VSEPR model of molecular
geometry. The BrO3F2- anion is only the fourth Br(VII) species
to have been isolated in macroscopic quantities and structurally
characterized and is one of only three oxide fluorides known
to possess D3h symmetry, the others being XeO3F2 and matrix-
isolated OsO3F2 monomer.
-
3.04 mdyne Å-1), and BrO3F2 (2.09, 2.19 mdyne Å-1), with
-
the order reversed at ClO3F2 (1.68, 1.73 mdyne Å-1). The
orderings cannot, however, be accounted for in terms of the fdd
contributions to F22 and F33 derived from experimental frequen-
cies. The higher symmetry force constants of the symmetric
-
The BrO3F2 salts described in this work provide experi-
mental confirmation of the fluoride ion affinity of BrO3F
-
and antisymmetric BrF2 stretches of BrO3F2 relative to those
calculated in this and in previous work.26 Attempts to prepare
of ClO3F2- are in accord with the greater fluoride ion basicity
-
2-
salts containing the ClO3F2 and BrO3F3 anions by the
reactions of ClO3F and [N(CH3)4][BrO3F2] with [N(CH3)4][F]
were unsuccessful. Gas-phase electronic structure calculations
confirm that the fac- and mer-isomers of the BrO3F32- dianion
-
of ClO3F2 and with the Br-F bond order of BrO3F2-, which
-
is 45% greater than the Cl-F bond order of ClO3F2 (vide
-
-
infra). Decreases in F44 along the series ClO3F2 > BrO3F2
> XeO3F2 > OsO3F2 reflect decreasing resistance to angular
distortions for the out-of-plane EO3 bend with increasing size
of the central atom, with F77 and F88 of the in-plane EO2 bends
showing similar, but less pronounced, trends.
-
-
are unstable, whereas ClO3F2 and BrO3F2 are predicted to
be thermodynamically stable with respect to fluoride ion loss.
The enthalpy of fluoride ion attachment for ClO3F (-132 kJ
mol-1) is, however, considerably less exothermic than that of
BrO3F (-261 kJ mol-1) and is comparable with the minimum
value typically required for the formation of a stable anion
(-126 to -146 kJ mol-1) when lattice and solvation energies
(e) Atomic Charges and Bond Orders. The atomic charges,
-
valencies, and bond orders of XO4-, XO3F, and XO3F2
(X ) Cl, Br) are provided in Table S6 and were calculated by
use of the natural bond orbital (NBO) method in conjunction
with the energy-minimized geometries determined by the LDF
(MPW1PW91) method.
-
are taken into account.26 The calculated BrO3F2- and ClO3F2
-
bond orders indicate that the X-O bonds of ClO3F2 are
stronger but the X-F bonds of ClO3F2- are weaker than those
9
J. AM. CHEM. SOC. VOL. 127, NO. 26, 2005 9425

A R T I C L E S
Lehmann and Schrobilgen
of BrO3F2-, which is consistent with the internal and symmetry
force constant analyses of these anions and the greater fluoride
ion affinity of BrO3F relative to that of ClO3F.
by monitoring the intensities of the strong RbF bands at 3144(100),
3160(87), 3338(4), 3357(2), and 3396(21) cm-1. Typically, CH3CN
solvent (0.25 mL) was condensed into a 4-mm or 1/4-in. o.d. FEP vessel
containing 0.1 to 0.2 mmol of powdered MF (M ) Cs, Rb, K, N(CH3)4).
The vessel was backfilled with N2 and temporarily stored at -78 °C
to prevent the alkali metal fluorides and, in particular, [N(CH3)4][F]37,38
from reacting with the solvent. A stoichiometric excess of BrO3F (ca.
0.6 mmol) was condensed into the vessel at -196 °C (vide supra), and
the vessel was backfilled with 1000 Torr of Ar at -78 °C. The mixture
was warmed to above the melting point of the CH3CN solvent for 1 h
prior to isolation of the colorless salts by removal of the solvent and
excess BrO3F under dynamic vacuum at this temperature. The optimum
reaction temperature varied from salt to salt, likely reflecting the relative
solubilities of the fluoride ion donors. Cesium fluoride readily reacted
with BrO3F in CH3CN at -40 to -48 °C to produce R-[Cs][BrO3F2],
but formed â-[Cs][BrO3F2] when the solution was warmed to -35 °C.
The high-temperature â-phase was also prepared by warming
R-[Cs][BrO3F2] to 0 °C in the absence of a solvent while monitoring
the phase transition by Raman spectroscopy. Complete conversion of
R-[Cs][BrO3F2] to â-[Cs][BrO3F2] by the latter method required ca.
35 h. The reactions of KF and RbF with BrO3F proceeded very slowly
below -40 °C, but within an hour when warmed to between -30 and
-35 °C. The reaction of [N(CH3)4][F] with BrO3F proceeded at -40
°C; however, warming these CH3CN solutions above 0 °C led to slow
reaction of BrO3F2- with CH3CN to produce NO2F (δ (19F), 396.6 ppm;
Experimental Section
Caution: Anhydrous HF must be handled using appropriate protec-
tive gear with immediate access to proper treatment procedures78-80 in
the event of contact with liquid HF, HF vapor, or HF-containing
solutions. Perbromyl fluoride, ClO3F, and the BrO3F2- salts are strong
oxidants and may react vigorously to explosively with water, organic
compounds, and other oxidizable materials. The syntheses of [M][BrO3F2]
(M ) K, Rb, Cs, N(CH3)4) are specifically cautioned, as solutions of
BrO3F and CH3CN have, on occasion, detonated. Small-scale (<100
mg) syntheses of the aforementioned salts are recommended.
(a) Apparatus and Materials. All volatile materials were handled
in vacuum lines constructed of stainless steel, nickel, and FEP
fluoroplastic, whereas nonvolatile materials were transferred in a drybox
as previously described.81 Acetonitrile (Caledon, HPLC grade) was
purified by the literature method82 and transferred under static vacuum
on a glass vacuum line. Potassium perbromate,7 AsF5,83 NOF,84 and
anhydrous [N(CH3)4][F]37 were prepared and purified according to
literature methods. Sodium fluoride (J. T. Baker Chemical Co., 99%)
and KF (J. T. Baker Chemical, 99%) were dried in a glass vessel under
dynamic vacuum at 250-300 °C for a minimum of 3 days and stored
in a drybox. Cesium fluoride (Aldrich, 99.9%) and RbF (ICN-K&K
Laboratories Inc., 99.9%) were dried by fusion in a platinum crucible,
and the molten salts were allowed to cool under dynamic vacuum in
the antechamber of a drybox. The resulting CsF and RbF pellets were
ground inside a drybox and stored in PFA or FEP vessels. Anhydrous
HF was purified as previously described39 and stored over BiF5 in a
Kel-F vessel. High-purity Ar or N2 gases were used for backfilling
vessels.
3
3J(19F-14N), 115 Hz), CH3COF (δ (19F), 49.0 ppm; J(19F-1H), 7.2
Hz), and BrO4 (δ (79Br), 2482 ppm).
-
[NO]2[BrO3F2][F]. The salt, [NO]2[BrO3F2][F], was synthesized by
the reaction of BrO3F with liquid NOF at -78 °C. Perbromyl fluoride
(ca. 0.6 mmol) was condensed into a 1/4-in. o.d. FEP vessel containing
0.1 mL of NOF (ca. 0.13 g, 2.7 mmol) which, upon warming to -78
°C, resulted in a colorless solid. The vessel was briefly (ca. 30 s)
evacuated at -78 °C until the last visible traces of the liquid reagents
had evaporated, but was not pumped on further because the product
has an appreciable dissociation vapor pressure and rapidly pumps off
at -78 °C.
Attempted Syntheses of [N(CH3)4]2[BrO3F3] and [N(CH3)4]-
[ClO3F2]. The synthesis of [N(CH3)4]2[BrO3F3] was attempted by
addition of [N(CH3)4][F] (6.98 mg, 74.9 µmol) to [N(CH3)4][BrO3F2]
(6.48 mg, 69.6 µmol) at -160 °C in a 4-mm o.d. FEP vessel.
Acetonitrile (0.1 mL) was then condensed into the vessel, and the
solution was warmed to 0 °C for 5 min to dissolve the reagents. The
Raman spectrum of the colorless solid remaining after solvent removal
under dynamic vacuum at -40 °C was a mixture of [N(CH3)4][BrO3F2]
and [N(CH3)4)][F],38 indicating that transfer of a second fluoride ion
to BrO3F to form [N(CH3)4]2[BrO3F3] had not occurred.
BrO3F. Perbromyl fluoride was prepared as previously described1,24
by the reaction of AsF5 with [K][BrO4] in anhydrous HF at ambient
temperature. In a typical reaction, 110 mg (0.601 mmol) of [K][BrO4]
1
was dissolved in 0.75 mL of HF in a /4-in. o.d. FEP vessel equipped
with a Kel-F valve. Arsenic pentafluoride (0.306 g, 1.80 mmol) was
then condensed into the vessel at -196 °C, and the vessel was backfilled
with Ar after dissolution of the AsF5 in HF at -78 °C. The vessel and
contents were warmed to ambient temperature for 45-60 min with
periodic mixing to ensure complete reaction. The solution was again
cooled to -78 °C, and the volatile components were distilled under
dynamic vacuum into a 5/8-in. o.d. FEP U-tube (-196 °C) containing
3 g of dry NaF. The U-tube and contents were isolated under static
vacuum and warmed to ambient temperature using a water bath, which
served as a heat sink for the exothermic reactions of HF and AsF5
with NaF. After 5-10 min, the U-tube was cooled to -196 °C, and
the vacuum in the tube was refreshed. Purified BrO3F was transferred
by rapidly warming the U-tube to ambient temperature using a water
bath and condensing the volatile product under static vacuum into a
1/4-in. or 4-mm o.d. FEP reaction vessel at -196 °C.
The synthesis of [N(CH3)4][ClO3F2] was attempted by analogy with
that of [N(CH3)4][BrO3F2]. Perchloryl fluoride (80.6 µmol, Pennsalt
Chemicals) was condensed into an FEP vessel (4-mm o.d.) containing
a preweighed amount of [N(CH3)4][F] (7.51 mg, 80.6 µmol) dissolved
in 0.3 mL of CH3CN. The vessel was then warmed to -40 °C for 1 h
with agitation before removal of the solvent at -40 °C under dynamic
vacuum. The Raman spectrum of the remaining colorless solid, recorded
at -40 °C, corresponded to that of [N(CH3)4][F].38 The synthesis of
[N(CH3)4][ClO3F2] was also attempted by the direct reaction of
[N(CH3)4][F] (28.5 mg; 0.306 mmol) with a large excess of ClO3F
(0.11 mL; 1.5 mmol) in a 4-mm o.d. FEP vessel. The reaction mixture
was warmed to -40 °C for 2 h with periodic agitation. The Raman
spectrum (-163 °C) of the colorless solid under ClO3F was consistent
with a mixture of [N(CH3)4][F] and ClO3F and provided no evidence
for [N(CH3)4][ClO3F2] formation. The Raman spectrum (-163 °C) of
the solid remaining after the removal of the ClO3F at -120 °C was
that of [N(CH3)4][F].
[M][BrO3F2] (M ) K, Rb, Cs, N(CH3)4). The title salts were
prepared by reaction of BrO3F with the alkali metal fluorides and
-
[N(CH3)4][F] in CH3CN. The formation of the BrO3F2 salt and
consumption of BrO3F were monitored by periodically freezing the
reaction mixtures and recording their Raman spectra at -163 °C. In
the case of [Rb][BrO3F2], the progress of the reaction was also followed
(77) Levin, I. W. J. Mol. Spectrosc. 1970, 33, 61.
(78) Bertolini, J. C. J. Emerg. Med. 1992, 10, 163.
(79) Peters, D.; Miethchen, R. J. Fluorine Chem. 1996, 79, 161.
(80) Segal, E. B. Chem. Health Saf. 2000, 19.
(81) Casteel, W. J., Jr.; Kolb, P.; LeBlond, N.; Mercier, H. P. A.; Schrobilgen,
G. J. Inorg. Chem. 1996, 35, 929.
(82) Winfield, J. M. J. Fluorine Chem. 1984, 25, 91.
(83) Mercier, H. P. A.; Sanders, J. C. P.; Schrobilgen, G. J.; Tsai, S. S. Inorg.
Chem. 1993, 32, 386.
(b) X-ray Crystallography. (i) Crystal Growth. Crystals of
[N(CH3)4][BrO3F2] were grown as previously described49 by slowly
cooling a CH3CN solution of the salt from 10 to -25 °C in a T-shaped
(84) Christe, K. O. Inorg. Chem. 1973, 12, 1580.
9
9426 J. AM. CHEM. SOC. VOL. 127, NO. 26, 2005

-
Synthesis of Salts Containing the BrO3F2 Anion
A R T I C L E S
1/4-in. o.d. FEP vessel. After crystal growth was complete, the solvent
was decanted into the sidearm of the T-reactor and cooled to -196
°C, and the sidearm was heat-sealed off and removed under dynamic
vacuum. The colorless, tetragonal-shaped crystals were dried under
dynamic vacuum at -40 °C and stored at -78 °C until mounted on
the diffractometer. Crystals of [NO]2[BrO3F2][F] were grown by
(e) Calculations. The energy-minimized gas-phase structures,
vibrational frequencies, atomic charges, Mayer bond orders, and
valencies were calculated by use of the HF, MP2, and LDF
(MPW1PW91) methods using Gaussian 98.90 The 6-311G(d) (HF, MP2)
and DZVP (MPW1PW91) basis sets were used in each case. The
-
enthalpies of fluoride ion attachment to ClO3F, BrO3F, and BrO3F2
1
sublimation of the initial product under 1000 Torr of Ar in a /4-in.
were calculated by use of the Gaussian-2 (G2) method.
The internal force constants of BrO3F2-, XeO3F2, and OsO3F2 were
FEP vessel over the course of several weeks while stored in a dewar
filled with solid dry ice. The colorless, block-shaped crystals ac-
cumulated in the cooler upper region of the vessel where fresh dry ice
had been added daily.
(ii) Crystal Mounting and Data Collection. A single crystal of
[N(CH3)4][BrO3F2] was mounted on a glass fiber at -110 ( 5 °C using
a Fomblin oil as an adhesive.85 A crystal of [NO]2[BrO3F2][F] was
mounted in a similar fashion; however, the substantial dissociation vapor
pressure of this salt and its tendency to sublime in a cold stream of dry
N2 at -110 °C required that it be mounted at -150 °C.
determined by the Wilson FG-matrix method74 using the software
package SVIB.91 The experimental geometry was used for BrO3F2
,
-
and the calculated LDF (MPW1PW91) values were used for XeO3F2
-
and OsO3F2 monomer. The symmetry force constants of BrO3F2
,
ClO3F2-, XeO3F2, and OsO3F2 were determined by use of the B-matrix
method using the second derivative potential energy matrixes obtained
from their calculated structures (SVWN/DZVP) and the software
package B-matrix.75
Acknowledgment. We thank the donors of the Petroleum
Research Fund, administered by the American Chemical Society,
for support of this work under ACS-PRF No. 37128-AC3
(G.J.S.), and the Natural Sciences and Engineering Research
Council of Canada for a postgraduate scholarship and McMaster
University for a Dalley Fellowship (J.F.L.). We are grateful to
Prof. David A. Dixon for his assistance with the use of the
B-matrix software package and to Dr. He´le`ne P. A. Mercier
for her considerable assistance in preparing and critiquing this
manuscript.
Mounted crystals were centered on a P4 Siemens diffractometer
(-173 °C) equipped with a Siemens SMART 1K CCD area detector,
a rotating molybdenum anode (λK R ) 0.71073 Å, monochromated by
a graphite crystal) and controlled by SMART.86 The distance between
the crystal and the detector face was 4.970 ([NO]2[BrO3F2][F]) or 4.987
cm ([N(CH3)4][BrO3F2]), and the data sets were collected in 512 ×
512 pixel mode using 2 × 2 pixel binning. The raw diffraction data
sets were integrated and scaled as previously described49 using
SAINT+87 and SADABS.88
(iii) Solution and Refinement. The program XPREP89 was used to
confirm unit cell dimensions and space groups. Direct methods were
used to locate the bromine atoms, and the lighter atom positions were
identified in successive difference Fourier syntheses. Final refinements
were obtained using data that had been corrected for absorption by
introducing an extinction coefficient and were optimized using aniso-
tropic thermal parameters for all atoms except the hydrogen atoms of
the N(CH3)4+ cation. Attempts to constrain the bond angles and/or bond
lengths of the anion in the disordered structure of [N(CH3)4][BrO3F2]
did not improve the global solution and therefore were not utilized in
the final refinement of the structure.
Supporting Information Available: Calculated vibrational
-
frequencies of BrO3F2 and ClO3F2-; stretching frequency
-
trends among BrO3F2 salts; factor-group analysis of the
-
BrO3F2 anion in [NO]2[BrO3F2][F]; calculated geometries of
fac-BrO3F32- and mer-BrO3F32- and enthalpies of fluoride ion
attachment to BrO3F2- (complementary discussion); calculated
-
vibrational frequencies and intensities for the BrO3F2 and
-
ClO3F2- anions (Table S1); factor-group analysis of the BrO3F2
(c) NMR Spectroscopy. The solution 19F (470.592 MHz), 1H
(500.130 MHz), and 13C (125.758 MHz) spectra of [N(CH3)4][BrO3F2]
in CH3CN solvent at -40 °C were recorded on a Bruker DRX-500
(11.7438 T) spectrometer operating in unlocked mode (field drift <
0.1 Hz h-1) using a 5-mm 1H/13C/19F/31P QNP probe. The spectra were
externally referenced to neat CFCl3 (19F) and Si(CH3)4 (1H, 13C) at 27
°C. The following acquisition parameters were used: pulse widths 2.3
(19F), 7.7 (1H), 12.8 (13C) µs; acquisition times 0.174 (19F), 2.42 (1H),
0.565 (13C) s; spectral widths 18.8 (19F), 6.8 (1H), 29.0 (13C) kHz. Free
induction decays were recorded in 32 K memories, zero filled to 64 K
memories and Fourier transformed, resulting in data point resolutions
of 2.9 (19F), 0.21 (1H), 0.88 (13C) Hz/data point.
anion in [NO]2[BrO3F2][F] (Table S2); calculated geometric
2-
parameters of BrO3F2-, ClO3F2-, and BrO3F3 (Table S3);
calculated gas-phase enthalpies of fluoride ion attachment (Table
S4); potential energy distributions for BrO3F2-, XeO3F2, and
OsO3F2 derived from GVFF analyses (Table S5); the disordered
structure of [N(CH3)4][BrO3F2] (Figure S1); displacement
-
vectors for the fundamental vibrational modes of the BrO3F2
anion (Figure S2); and X-ray crystallographic files (CIF format)
for the structure determinations of [NO]2[BrO3F2][F] and
[N(CH3)4][BrO3F2]. This material is available free of charge
via the Internet at http://pubs.acs.org.
(d) Vibrational Spectroscopy. Raman spectra were recorded on a
Bruker RFS 100 FT-Raman spectrometer at -163 °C using 1064-nm
excitation as previously described.85 The infrared spectrum of
â-[Cs][BrO3F2] was recorded on a Bio-Rad FTS-40 spectrometer at
ambient temperature. The three-layered AgCl pellet of [Cs][BrO3F2]
was fabricated in a drybox by use of a Wilks mini-press, with the
external layers being composed of AgCl and the central layer being a
mixture of the sample and AgCl.
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Instrumentation: Madison, WI, 1999.
(87) SAINT+, version 6.02; Siemens Energy and Automotive Analytical
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