Vibrational Spectra and Structural Aspects of Fluorosulfites

Article abstract of DOI:10.1021/ic970315o

The Raman and infrared spectra of the fluorosulfites of K⁺, Rb⁺, Cs⁺, NH₄⁺, and Me₄N⁺ have been examined. Previous assignment of the fundamental vibrations is revised, and an ab initio study of the SO₂F^- anion is presented. For the pyramidal anion of symmetry C_s, distances of r(S-O) = 1.458 ? and r(S-F) = 1.698 ? have been calculated. The heat of formation of Me₄NSO₂F (-14.0 kcal/mol) was derived from the dissociation pressure of the salt. In addition, the lattice energy of Me₄NF (159.2 kcal/mol) was calculated from a thermochemical cycle. The thermochemical data are discussed in terms of reactivity of the fluorides and stability of the fluorosulfites.

Full text of DOI:10.1021/ic970315o

5570
Inorg. Chem. 1997, 36, 5570-5573
Vibrational Spectra and Structural Aspects of Fluorosulfites
Andreas Kornath* and Frank Neumann
Anorganische Chemie, Fachbereich Chemie, Universita¨t Dortmund, 44221 Dortmund, Germany
Ralf Ludwig
Physikalische Chemie, Fachbereich Chemie, Universita¨t Dortmund, 44221 Dortmund, Germany
ReceiVed March 18, 1997X
The Raman and infrared spectra of the fluorosulfites of K+, Rb+, Cs+, NH4+, and Me4N+ have been examined.
Previous assignment of the fundamental vibrations is revised, and an ab initio study of the SO2F- anion is presented.
For the pyramidal anion of symmetry Cs, distances of r(S-O) ) 1.458 Å and r(S-F) ) 1.698 Å have been
calculated. The heat of formation of Me4NSO2F (-14.0 kcal/mol) was derived from the dissociation pressure of
the salt. In addition, the lattice energy of Me4NF (159.2 kcal/mol) was calculated from a thermochemical cycle.
The thermochemical data are discussed in terms of reactivity of the fluorides and stability of the fluorosulfites.
Introduction
we found in the course of our investigations of alkali metal
fluoride catalyzed reactions that fluorosulfites, which are readily
formed in the presence of traces of sulfur dioxide, act as
deactivators. This apparent contradiction and a recently found
unusual reaction of CsF with SO2 enhanced by ultrasound (eq
1)18 gave rise to investigations of the fluorosulfinate salts.
Fluorosulfites were mentioned for the first time by Seel and
Meier in the course of their investigation of the lead chamber
process.1 Later, Seel et al. described the preparation of
Me4N+SO2F- by the reaction of sulfur dioxide with naked
fluoride (Me4NF) and reported the first studies of the Na+, K+,
Rb+, Cs+, and Me4N+ salts by X-ray powder patterns.2,3 In
the case of the potassium salt, similarities to KClO3 were found,
but no structural data were obtained. The examination of the
dissociation pressure of fluorosulfites was followed by the first
reported IR spectrum of a fluorosulfite salt (KSO2F).4-6
Robinson et al. reported the IR spectra of the K+, Rb+, Cs+,
and NMe4+ salts and calculated force constants for the fluoro-
sulfite ion.7 The results seem to be consistent with the work
of Seel and Boudier,6 but the spectra exhibit three additional
bands in the region between 600 and 1100 cm-1 and one at
1230 cm-1, all attributed to combination tones. However, the
next spectroscopic studies were reported 22 years later by
Moock et al., who described the preparation of NH4SO2F.8
Despite the fact that fluorosulfite is the smallest fluorooxysulfur
anion, no work has been done to obtain the geometric parameters
by theoretical investigations.
2CsF + 4SO2 f Cs2S3O6 + SO2F2
(1)
Furthermore, the reversible reaction of fluorides with SO2
enables us to study the differences in the series KF, RbF, CsF,
and Me4NF, taking thermochemical parameters into account.
Experimental Section
Apparatus and Materials. All synthetic work and sample handling
were performed by employing standard Schlenk techniques and a
standard vacuum line. The alkali metal fluorides were dried thoroughly
and crushed in a drybox under nitrogen. NH4F was prepared from dry
HF19 and NH3, since attempts to purify commercially available NH4F
by the method of Moock et al. gave a mixture of NH4F and NH4HF2.8
Me4NF was obtained by the known literature method.20 SO2 (Messer
Griesheim) was dried over CaH2. (Caution! Anhydrous HF causes
skin burning and irreparable damages! Fluorides and fluorosulfites
are toxic!)
The Raman spectra were recorded on a T64000 (ISA) using an argon
ion laser (Spectra Physics) at 514.5 nm. The infrared spectrometer
and sample preparation were described previously.19 The elemental
analysis of fluoride and sulfur (as SO42-) was carried out by ion
chromotography of samples dissolved in water and oxidized with iodine.
Preparation of the Fluorosulfites. A 500 mg sample of carefully
dried alkali metal fluoride (KF, RbF, CsF) was placed in a dried glass
tube, and 5 g of SO2 was condensed into the tube at -196 °C. The
sealed glass tube was warmed to room temperature. The reaction
proceeded upon thawing in a few minutes. The mixture was stirred
This lack of structural parameters for the SO2F- anion was
surprising since KSO2F was proposed as actiVated fluoride for
a long time.9,10 In fact, besides the investigations by Seel et
al.,2,3 it was used for only a few syntheses.11-17 Furthermore,
X Abstract published in AdVance ACS Abstracts, October 15, 1997.
(1) Seel, F.; Meier, H. Z. Anorg. Allg. Chem. 1953, 274, 202.
(2) Seel, F.; Jonas, H.; Riehl, L.; Langer, J. Angew. Chem. 1955, 67, 32.
(3) Seel, F.; Riehl, L. Z. Anorg. Allg. Chem. 1955, 282, 293.
(4) Seel, F.; Go¨litz, D. Z. Anorg. Allg. Chem. 1964, 327, 28.
(5) Paetzold, R.; Aurich, K. Z. Anorg. Allg. Chem. 1965, 335, 281.
(6) Seel, F.; Boudier, J. Z. Anorg. Allg. Chem. 1966, 342, 173.
(7) Robinson, E. A.; Lavery, D. S.; Weller, S. Spectrochim. Acta 1968,
25A, 151.
(14) Bamgboye, T. T.; Bamgboye, O. A. Spectrochim. Acta 1984, 40A,
329.
(8) Moock, K. H.; Su¨lzle, D.; Klaeboe, P. J. Fluorine Chem. 1990, 47,
151.
(9) Seel, F. Inorg. Synth. 1967, 9, 113.
(10) Kwasnik, W. In Handbuch der Pra¨paratiVen Anorganischen Chemie;
Brauer, G., Ed.; Enke Verlag: Stuttgart, Germany, 1975; p 244.
(11) Biran, Z.; Goldschmidt, J. M. E. Synth. React. Inorg. Met.-Org. Chem.
1978, 8, 323.
(12) Allenstein, E.; Schrempf, G. Z. Anorg. Allg. Chem. 1981, 474, 7.
(13) Bamgboye, T. T.; Sowerby, D. B. J. Inorg. Nucl. Chem. 1981, 43,
2253.
(15) Meisel, M.; Donath, C.; Cernik, M.; Wolf, G.-U.; Dostal, K.; Grunze,
H. Z. Anorg. Allg. Chem. 1984, 512, 72.
(16) Bamgboye, T. T.; Bamgboye, O. A. Spectrochim. Acta 1985, 41A,
981.
(17) Volkmann, R. A. PCT Int. Appl. WO 88 08,845, Nov 17, 1988.
(18) Kornath, A.; Neumann, F. Inorg. Chem., submitted for publication.
(19) Minkwitz, R.; Kornath, A.; Sawodny, W. Angew. Chem. 1992, 104,
648; Angew. Chem., Int. Ed. Engl. 1992, 31, 643.
(20) Christe, K. O.; Wilson, W. W.; Wilson, R. D.; Bau, R.; Feng, J. J.
Am. Chem. Soc. 1990, 112, 7619.
S0020-1669(97)00315-7 CCC: $14.00 © 1997 American Chemical Society

Vibrational Spectra of Fluorosulfites
Inorganic Chemistry, Vol. 36, No. 24, 1997 5571
for 1 day to ensure complete reaction. Excess SO2 was pumped off at
dry ice temperature. The increase in weight and the elemental analyses
of the white solids indicated quantitative yields in the cases of RbSO2F
and CsSO2F, but ca. 85% yield for KSO2F. Quantitative yields of
KSO2F were obtained using KF that was prepared by thermal
decomposition of the 85% KSO2F at 140 °C in high vacuum. Further
characterization of the obtained compounds was carried out by
vibrational spectroscopy.
NH4SO2F and Me4NSO2F were prepared by the same procedure as
CsSO2F, but the reaction mixture was warmed to -40 °C. The yields
were quantitative. The fluorosulfites hydrolyze rapidly upon contact
with moisture.
Dissociation Pressure of Me4NSO2F. Me4NSO2F decomposes
slowly under static SO2 pressure at temperatures around 100 °C. The
decomposition proceeds under formation of Me4NSO3F, Me4NF, and
sulfur. Besides SO2, small traces of Me3N, CH3F, and SO2F2 were
found in the gas phase. At temperatures below 100 °C only dissociation
of Me4NSO2F under formation of the starting materials occurs. The
SO2 pressure of Me4NSO2F was measured in the range of 1-400 hPa
(20-120 °C) using a calibrated pressure gauge (MKS).
Dissociation pressure (p (hPa)/T (°C)): 1.3/20, 2.5/30, 5.4/40, 10.6/
50, 19.3/60, 35.0/70, 62.0/80, 106/90, 172/100, 261/110, 400/120.
A logarithmic plot of the pressure versus the reciprocal temperature
yields a linear correlation (in the range of 20-100 °C) with the equation
log(p/hPa) ) 10.16-2950 K/T.
Attempts to obtain a reliable p/T correlation for NH4SO2F failed,
since the decomposition was accompanied by sublimation.
Computational Methods. The ab initio calculations for SO2F- and
SOF2 were performed at the restricted Hartree-Fock level of theory
using the Gaussian 94 program.21 All calculations were carried out at
the 6-31+G* basis level, which augments the standard double-ú plus
polarisation treatment (6-31-G*) with a diffuse set of s,p functions (+)
on each heavy atom and is known to describe anionic systems in an
appropriate way. Harmonic vibrational frequencies were computed for
the minimum-energy structures and scaled by the empirical factor 0.90
to maximize their fit with the experimentally observed frequencies.22,23
Figure 1. Infrared spectrum of KSO2F (trace A) and Raman spectra
of KSO2F (trace B), RbSO2F (trace C), and CsSO2F (trace D).
pressures up to ca. 200 hPa and can be used to obtain highly
reactive fluorides.18,24 At higher pressures of SO2, the salts
decompose under formation of the fluorosulfate and sulfur
according to eq 4. This disproportionation is accompanied by
Results and Discussion
Synthesis and Properties of Fluorosulfites. The salts of
Rb+, Cs+, NH4+, and Me4N+ are formed in quantitative yields
by reacting an excess of SO2 with the fluorides according to eq
2. In case of KF, only yields of about 85% were obtained.
3M+SO2F- f 2M+SO3F- + MF + 1/8S8
(4)
MF + SO2 f MSO2F (M ) K, Rb, Cs, NH4, Me4N) (2)
the decomposition of Me4NF, in the case of the Me4NSO2F (eq
5). The enthalpy of formation for this salt was derived from
Quantitative yields of KSO2F were achieved using KF that was
obtained by thermal decomposition of the 85% KSO2F18 or using
dimethyl sulfoxide as solvent.9
Me4NF f Me3N + MeF
(5)
The fluorosulfite salts are colorless microcrystalline salts that
rapidly decompose on contact with moisture according to eq 3.
the SO2 pressure at different temperatures and is discussed later.
The alkali metal and ammonium fluorosulfites are insoluble
in nonpolar and polar solvents such as SO2, CH3CN, and
(CH3)2SO. Only Me4NSO2F is soluble in SO2. In acetonitrile,
a fast decomposition and evolution of SO2 were observed. The
solution of Me4NSO2F in SO2 strongly tends to oversaturation.
Therefore it was not possible to obtain single crystals for X-ray
diffraction studies either by slow cooling of the solution or by
pumping off the solvent.
4M+SO2F- + H2O f 2M+HF2- + M2S2O5 + 2SO2 (3)
The thermal decomposition of the fluorosulfites is more
complex. The first step is a simple dissociation in SO2 and the
fluoride. This thermal decomposition is reversible for SO2
Vibrational Spectra. The Raman spectra of KSO2F, RbSO2F,
and CsSO2F together with the infrared spectrum of KSO2F are
shown in Figure 1, and the observed frequencies are summarized
in Table 1. The vibrational data for NH4SO2F and Me4NSO2F
are listed in Tables 2 and 3, respectively. The assignments for
SO2F- were made by comparison with the isoelectronic ClO2F
(Table 4).25 An ab initio investigation which is discussed later
(21) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Gill, P. M. W.; Johnson,
B. G.; Robb, M. A.; Cheeseman, J. R.; Keith, T. A.; Petersson, G. A.;
Montgomery, J. A.; Raghavachari, K.; Al-Laham, M. A.; Zakrzewski,
V. G.; Ortiz, J. V.; Foresman, J. B.; Cioslowski, J; Stefanov, B. B.;
Nanayakkara, A.; Challacombe, M.; Peng, C. Y.; Ayala, P. Y.; Chen,
W.; Wong, M. W.; Andres, J. L.; Replogle, E. S.; Gomperts, R.;
Martin, R. L.; Fox, D. J.; Binkley, J. S.; Defrees, D. J.; Baker, J.;
Stewart, J. P.; Head-Gordon, M; Gonzalez, C.; Pople, J. A. Gaussian
94, Revision A.1; Gaussian, Inc.: Pittsburgh, PA, 1995.
(22) Pulay, P.; Forgasi, G.; Pang, F.; Boggs, J. E. J. Am. Chem. Soc. 1979,
101, 2550.
(24) Kwasnik, W. In Handbuch der Pra¨paratiVen Anorganischen Chemie;
Brauer, G., Ed.; Enke Verlag: Stuttgart, Germany, 1975; p 244.
(25) Smith, D. F.; Begun, G. M.; Fletcher, W. H. Spectrochim. Acta 1964,
20, 1763.
(23) Pople, J. A.; Schlegel, H. B.; Krishnan, R.; Frees, D. J.; Binkley, J.
S.; Frisch, M. J.; Whiteside, R. A.; Hout, R. F.; Hehre, W. J. Int. J.
Quantum Chem. 1981, 15, 269.

5572 Inorganic Chemistry, Vol. 36, No. 24, 1997
Kornath et al.
assignment
Table 1. Vibrational Frequencies (cm-1) of the Alkali Metal Fluorosulfites
KSO2F
IRa
RbSO2F
IRa
CsSO2F
IRa
IR
Ra
IR
Ra
IR
Ra
90 (34)
103 (33)
248 (1)
370 (31)
388 (42)
88 (22)
122 (17)
247 (3)
370 (36)
387 (43)
112 (5)
131 (4)
243 (1)
lattice
lattice
lattice
ν6 (A′′) δas(OSF)
239 w
360 m
380 m
395 sh
498 s
587 m
1107 m
1171 s
1186 s
366 m
381 m
270
375
369 m
379 s
265
375
275
353
367 (10)
378 (12)
393 (15)
500 (12)
594 (5)
1105 (100)
1169 (5)
1183 (10)
ν4 (A′) δs(OSF)
}
504 s
596 s
1099 m
495
1105
1175
500 (18)
597 (9)
1109 (100)
499 s
595 s
1103 m
497
590
1102
497 (19)
594 (5)
1110 (100)
485
585
1090
ν3 (A′) δ(OSO)
ν2 (A′) ν(SF)
ν1 (A′) νs(SO)
ν5 (A′′) νas(SO)
}
1179 s
1187 (20)
1183 s
1180
1189 (17)
1163
a From ref 6.
Table 2. Vibrational Frequencies (cm-1) of Tetramethylammonium
Fluorosulfite
Table 4. Comparison of Observed and Calculated Frequencies
(cm-1) of SO2F-
NMe4+SO2F-
IRa
assignment
SO2F- (Cs)
obsd freqa
calcd
+
+
K+ Rb+ Cs+ NH4 NMe4 freqb ClO2Fc
assignment
IR
Ra
NMe4+ (Td)
368 369 363 371
384 383 387 416
502 498 497 518
596 594 590 588
1104 1106 1108 1100 1104 1086 1106 ν1 (A′) νs(XO)
1183 1186 1184 1167 1191 1174 1271 ν5 (A′′) νas(XO)
361
383
496
591
362
394
536
605
367 ν6 (A′′) δas(OXF)
402 ν4 (A′) δs(OXF)
546 ν3 (A′) δ(OXO)
630 ν2 (A′) ν(SF)
359 w
383 w
456 m
498 m
590 m
260
350
363 (16)
383 (7)
463 (16)
494 (2)
592 (5)
759 (31)
952 (31)
1105 (23)
1196 (4)
1295 (4)
ν6 (A′′) δas(OSF)
ν4 (A′) δs(OSF)
ν8 (E)
ν19 (F2)
495
596
767
950
1101
1187
ν3 (A′) δ(OSO)
ν2 (A′) ν(SF)
ν3 (A1)
ν18 (F2)
952 s
1103 s
1187 s
a Average values from infrared and Raman spectra. b Scaled by an
empirical factor of 0.9. c From ref 25.
ν1 (A′) νs(SO)
ν5 (A′′) νas(SO)
ν7 (E)
17 (F2)
ν6 (E)
16 (F2)
ν2 (A1)
ν
+
the Me4N+ and NH4 cations were made according to well
1418 w
1412 (5)
known literature data.20,26,27
1466 (21)
1474 (44)
1487 (4)
ν
The two oxygen-sulfur stretching modes of the SO2F- anion
are observed at 1180 and 1105 cm-1, respectively. The
additional bands in this region measured by Robinson et al.7
were observed after exposing the sample to moisture and can
therefore be attributed to the hydrolysis products according to
eq 3.
1494 s
ν15 (F2)
2822 (9)
2916 (20)
2964 (67)
3041 (100)
ν
14 (F2)
ν5 (E)
3029 w
3406 w
3468 w
3550 w
The sulfur-fluorine stretching mode occurs in the region of
600 cm-1 at somewhat lower wavenumbers than for the
isoelectronic ClO2F (630 cm-1). It is remarkably low in
a From ref 7.
comparison to those of the SOF2 molecule (799 and 737 cm-1 28
)
Table 3. Vibrational Frequencies (cm-1) of Ammonium
Fluorosulfite
and indicates a weak sulfur-fluorine bond in agreement with
the ab initio calculations.
NH4+SO2F-
IRa
assignment
SO2F- (Cs)
The three deformation modes are observed in the low-
wavenumber region below 600 cm-1. The OSO bending mode
occurs at around 500 cm-1. In contrast to previous data,6-8
the two OSF bending modes were found close together around
IR
Ra
Raa NH4+ (Td)
110 (14)
133 (1)
163 (4)
200 (8)
240 (14)
372 (15)
412 (15)
522 (4)
597 (7)
384 and 368 cm-1
.
The spectra of KSO2F and RbSO2F show similar band shapes
for the SO2F- anion, whereas CsSO2F exhibits a band splitting
for νas(SO) and δs(OSF). The average frequencies for each
fluorosulfite salt together with those of the isoelectronic ClO2F
and calculated frequencies are given in Table 4. The observed
frequencies for the fluorosulfite salts are almost constant, except
for ν3, ν4, and ν5 of the NH4+ salt. The deviations are probably
due to strong interactions between the anion and the protons of
the cation. Overall, the SO2F- frequencies show the expected
trend to appear at somewhat lower wavenumbers than those of
the isoelectronic ClO2F.
248
370
248
370
370 w
420 m
515 m
580 m
ν6 (A′′) δs(OSF)
ν4 (A′) δs(OSF)
ν3 (A′) δ(OSO)
ν2 (A′) ν(SF)
ν1 (A′) νs(SO)
ν5 (A′′) νas(SO)
483
580
450
600
1098 m 1114 1101 (100) 1098
1167 s
1167
1168 (6) 1150
1183 (5)
1404 (10) 1403
1712 (3)
1759 (2)
3097 (52) 3087
1397 vs 1399
ν4 (F2)
ν2 (E)
ν2 (E)
ν1 (A1)
ν2 + ν4
ν3 (F2)
3094 sh
3216 s
Ab Initio Calculation. The ab initio RHF/6-31+G* calcu-
lated harmonic frequencies for the SO2F- anion are given in
Table 4 and compared with observed frequencies. We find
3184 (29)
a From ref 8.
(26) Christe, K. O.; Dixon, D. A.; Mahjoub, A. R.; Mercier, H. P. A.;
Sanders, J. C. P.; Seppelt, K.; Schrobilgen, G. J.; Wilson, W. W. J.
Am. Chem. Soc. 1993, 115, 2696.
(27) Mathieu, J. P.; Poulet, H. Spectrochim. Acta 1960, 16, 696.
(28) Kornath, A; Hartfeld, N. J. Mol. Spectrosc., in press.
gives, in accordance with the isoelectronic ClO2F, the expected
structure of symmetry Cs. Consequently, six fundamentals (4A′
+ 2A′′) which should all be active in both the infrared and the
Raman spectra are expected for SO2F-. The assignments of

Vibrational Spectra of Fluorosulfites
Inorganic Chemistry, Vol. 36, No. 24, 1997 5573
Table 5. Geometry of SO2F- (Calculated) Compared to Those of ClO2F, SOF2, and SO3F-
SOF2
SO2F-
calc
ClO2F
expb
SO3F-
expc
calc
expa
r(X-F) (Å)
1.698
1.458
113.2
100.6
1.576
1.409
1.583(3)
1.420(3)
1.694(3)
1.420(3)
115.2(5)
101.8(1)
1.555(6)
r(X-O) (Å)
1.424(4); 1.455(6)
113.5(2); 117.4(4)
104.5(5); 102.8(3)
(O-X-O) (deg)
(O-X-F) (deg)
(F-X-F) (deg)
106.6
92.5
106.2(2)
92.2(2)
a From ref 29. b From ref 30. c From ref 31.
Table 6. Thermochemical Data of the Fluorosulfites
dec pt
∆H°
(kcal/mol)
∆S°
(cal/(mol K))
∆U
(kcal/mol)
U(MF)
(kcal/mol)
U(MSO2F)
(kcal/mol)
(°C)a
KSO2F
203b
267b
329b
139
-18.3b
-20.8b
-23.2b
-14.0
38.4b
38.4b
38.4b
34.1
25.5
23.0
20.6
29.8
189.7c
181.6c
173.7c
159.2
164.2
158.6
153.1
129.4
RbSO2F
CsSO2F
Me4NSO2F
a Extrapolated for p ) 1013 hPa. b From ref 6. c From ref 35.
reasonable agreement for all vibrational modes. In particular,
the ab initio calculations support the experimental result that
SO2F- frequencies appear throughout at somewhat lower
wavenumbers than those of the isoelectronic ClO2F.
this apparent contradiction better, a comparison of the lattice
energies was made.
Lattice Energy of Me4NF. The lattice energies (U) are well-
known for the alkali metal halides,35 but no data are available
on Me4NF. To derive the lattice energy of Me4NF, we chose
the thermochemical cycle described by Wilson (eq 7).36
The calculated geometry for SO2F- and the isoelectronic
SOF2 are presented in Table 5. For SOF2 we find an almost
perfect agreement with experimental geometries measured by
Hargittai et al.29 Therefore, we are confident to predict correct
geometries for the SO2F- anion. The pyramidal SO2F- anion
of symmetry Cs has bond angles slightly different from the
isoelectronic ClO2F.30 The S-O bond length of 1.458 Å is
somewhat longer than the S-O bond in the isoelectronic SOF2
and in SO3F- 29,30 but still in the range of regular sulfur-oxygen
double bonds. The obtained S-F bond length of 1.698 Å is
remarkably long in comparison to those of SOF2 and SO3F- 31
but similar to the weak Cl-F length bond in the isoelectronic
ClO2F. Furthermore the weak S-F bond is in accordance with
the thermochemical data of the fluorosulfites.
U ) ∆Hf°(Me4N+, g) + ∆Hf°(F-, g) - ∆Hf°(Me4NF, c) +
Σ∆H + 298[Cp(Me4NF, c) - Cp(Me4N+, g) -
∫
T
0
Cp(F-, g)] dT (7)
The known thermochemical data were used: ∆Hf°(Me4N+,
g) ) 129.8 kcal/mol,33 ∆Hf°(F-, g) ) -59.9 kcal/mol,37
∆Hf°(Me4NF, c) ) -89.0 kcal/mol.38 The term Σ∆HT contains
any polymorphic transitions which Me4NF may undergo be-
tween 0 and 298 K. Such transitions are usually of the order
0.25 kcal/mol or less and have been ignored for the calculation.39
The heat capacity term is in the order of 0.28 kcal/mol when
Enthalpies of Formation. As already discussed, the reaction
of fluorides with SO2 is reversible under certain conditions. Seel
et al. derived from the p/T relation the enthalpy (∆Hf°) and
entropy (∆Gf°) of formation for the KSO2F, RbSO2F, and
CsSO2F.6 These values and the data obtained by us for
Me4NSO2F are summarized in Table 6. ∆Gf° for the Me4N+
salt is in the range for the alkali metal salts, whereas the ∆Hf°
of -14.0 kcal/mol is surprisingly low. We expected a higher
value, since Me4NF is known as the most reactive fluoride in
this series and the Me4N+ cation is known to stabilize anions
like PF4- and CO2F-,32,33 which are not available as alkali metal
salts.
298
the values of ∫0 Cp dT described by Wilson are used.36 The
equation yields a lattice energy of 159.2 kcal/mol. Although
this calculation incorporates some approximations, it should be
more justified than extended-calculation procedures which yield
differences for other Me4N+ salts in the order of 15 kcal/mol.
Such differences are explained by nonionic contributions to the
bonding.
The calculated lattice energy of Me4NF is about 14 kcal/mol
smaller than that for CsF. Therefore, reactions with Me4NF
require a lower activation energy and Me4NF exhibits a higher
reactivity than CsF. The fact that the lattice energy decreases
more rapidly in the case of the formation of Me4NSO2F than
for CsSO2F may be specific for this reaction and cannot be
generalized, since thermochemical data are rare in this field.
Since the gas phase F- affinity of SO2 is known (-43.8 kcal/
mol), a rough estimation of the change in lattice energy (∆U)
can be made by eq 6 derived from the Born-Haber cycle,
∆U ) ∆Hf° + 43.8 kcal/mol
(6)
ignoring differences in heat capacity and polymorphic transitions
between 0 and 298 K.34 The ∆U values in Table 6 indicate the
largest decrease in lattice energy for Me4NSO2F. To understand
Acknowledgment. This work was supported by the Deutsche
Forschungsgemeinschaft. Grateful thanks are expressed to Prof.
Rolf Minkwitz for helpful discussions.
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