(Hons.) Mcgill University, Montreal, 1966
OXYFLUORIDES, OXYFLUOROSULPHATES AND OXYCATIONS OF
THE HALOGENS AND SELENIUM
by
HENRY ALBERT CARTER
B.Sc. (Hons.) McGill University, Montreal, 1966
A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE
REQUIREMENTS FOR THE DEGREE OF
DOCTOR OF PHILOSOPHY
in the Department
of
CHEMISTRY
We accept this thesis as conforming to the
required standard
THE UNIVERSITY OF BRITISH COLUMBIA In presenting this thesis in partial fulfilment of the requirements for an advanced degree at the University of British Columbia, I agree that the Library shall make it freely available for reference and study.
I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the Head of my Department or by his representatives. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission.
Department of C k €. VA <-S"f f ^
The University of British Columbia Vancouver 8, Canada
Date December /
ABSTRACT
Alternate routes for the preparation of chloryl fluorosulphate,
ClO^SO^Fj have been found. Evidence for the existence of the chloronium
+ cation, C102 , as a red-coloured species in fluorosulphuric acid, HSO^F, has been obtained from electrical conductivity measurements, employing
S dichlor 1 C10"2S0 F and (C10"2)2 3°io' y trisulphate, as solutes.
nuclear magnetic resonance and uv-visible spectroscopy were also used to study solutions of ClC^SOgF in HSO^F, and are in agreement with the conductivity results.
The infrared and Raman spectra of the white solids, chloryl hexafluoroarsenate, chloryl hexafluoroantimonate and dichloryl hexafluorostannate have been obtained. The observed vibrations have been interpreted in terms of ionic compounds where strong cation-anion interaction via fluorine bridging results in a symmetry lowering of the anion. The effect of this interaction is also seen in the
Mossbauer spectrum of (ClO^^nF^, where a non-zero quadrupole splitting is found. All compounds dissolve in HSO^F to give red-coloured solutions
+ as a result of the formation of C1C>2 cations. Tn addition, conductivity
measurements show chloryl fluoride, C102F, to be completely'.dissociated
+ in HSO^F and to form C102 cations and SO^F" anions.',
The vibrational spectra of ClF2AsF6, ClF2SbF6 and BrF^bF^ have been recorded. In all compounds, various degrees of anion-cation interaction are found, resulting in the lowering of the symmetry of the anion. In addition, strong coupling of the vibrational modes for BrF„SbF, are observed. ii
An alternate route to the preparation of KBrfSO^F)^ was found.
A Raman spectrum of this compound indicates a square planar configuration
for the anion with polar bromine-oxygen bonds. The Raman spectra obtained for the tris-fluorosulphates ICSO^F)^ and Br(SO^F)^ indicates
the presence of two types of fluorosulphates groups and polymeric
structures for these compounds.
The structure of iodyl fluorosulphate, IC^SO^F, has been examined by Raman spectroscopy and conductivity measurements in HSO^F.
Evidence is found for discrete I02 groups linked by covalent bidentate fluorosulphate groups. A Raman spectrum of iodyl fluoride
showed it to be dissimilar to the chlorine analogue, CIC^F, and
possess a polymeric structure. No discrete I02 groups are found in
IO^AsFg on the basis of the infrared spectrum.
The new compound selenium (IV) oxyfluorosulphate, SeOCSO^F),,,
w t 1 has been prepared by the reaction of $2®^2 -*- ' SeOCl2 or SeO,,.
The structure of SeO(S03F)2 has been investigated mainly by Raman
spectroscopy and conductimetry in HS03F. A self-dissociation process
+ of the type,SeOCS03F)2 = SeOCS03F) + S03F~,is suggested for
SeO(S03F)2 to explain the results of the physical methods. Also,
the addition of SeOCl2 to SeO(S03F)2 results in the formation of
SeOCSO_F)C1. Both SeO(SO F)C1 and SeO(SO F) show unusually high
Se=0 vibrational stretching frequencies indicating polar Se-0S0„F bonds. iii
TABLE OF CONTENTS Page Abstract -'-i
Table of Contents iii
List of Tables vi
List of Figures V11
Acknowledgements ix
CHAPTER
1. INTRODUCTION
1.1 Oxyfluorides of the Halogens and Selenium 1
1.2 Scope of the Present Research 4
1.3 Preparative Routes to Fluorosulphates 5
1.4 Preparative Routes to Polyfluoroanion Complexes 7
2. GENERAL EXPERIMENTAL TECHNIQUES
2.1 Vacuum Systems and Apparatus 9
2.2 Reagents 20
2.3 Physical Experimental Methods 21
3. CHLORYL FLUOROSULPHATE: CHLORYL COMPOUNDS (RED) PART I
3.1 Introduction 27
3.2 Experimental 28
3.3 Results and Discussion 37
4. CHLORYL HEXAFLUOROMETALLATES: CHLORYL COMPOUNDS (WHITE) PART II
4.1 Introduction 60
4.2 Experimental 61
4.3 Results and Discussion 62 iv
5. DICHLORYL HEXAFLUOROSTANNATE:
CHLORYL COMPOUNDS (WHITE) PART II CONTINUED
5.1 Introduction 79
5.2 Experimental 80
5.3 Results and Discussion 81
6. STRUCTURAL STUDIES OF HEXAFLUOROARSENATES AND -ANTIMONATES
OF FLUORO-HALOGEN.HETEROCATIONS
6.1 Introduction 95
6.2 Experimental 99
6.3 Results and Discussion 99
7. BROMYL COMPOUNDS
7.1 Introduction 117
7.2 Experimental 118
7.3 Discussion 119
8. RAMAN STUDIES OF Br(S03F)3, I(S03F)3 AND THE ANIONS
Br(S03F)4" AND I(S03F)4~
8.1 Introduction 122
8.2 Experimental 123
8.3 Discussion 124
9. IODYL COMPOUNDS
9.1 Introduction 134
9.2 Experimental 136
9.3 Results and Discussion 141
10. SELENIUM (IV) OXYFLUOROSULPHATE
10.1 Introduction 160
10.2 Experimental 160
10.3 Results and Discussion 164 V
11. CONCLUSIONS AND SUMMARY 191
REFERENCES 194
APPENDIX 204 vi
LIST OF TABLES Page
1. Oxyfluorides of the Halogens and Selenium 2
2. Electrical Conductivity of C102S03F 40 3. First Ionization Potentials of Diatomic and Triatomic Molecules 42
4. Specific Conductivities of ClO^O^, (C102)2S3010
and K S 0 in HS03F 44
C10 S in H S0 49 5. Specific Conductivities of ( 2)2 3°10 2 4
1 6. H NMR Chemical Shifts of Solutions of C102S03F in HS03F 53
7. Electronic Spectra of C102S03F in HSC^F 55 8. The Vibrational Spectra of C10„AsF. and C10„SbF. 63 ZD Z O
9. Assignment of the C102 Vibrations 64
10. C102- Stretching Frequencies 64 11. Assignment of the AsF^ Vibrations 69 12. Assignment of the SbF^ Vibrations 69
13. Specific Conductivities of CIO,,- compounds in HS03F 74 14. Vibrational Frequencies for SnF^~ Compounds 82 15. Vibrational Frequencies for the SnF^ Anion 83
+ 16. Vibrational Frequencies of C102 in (ClO^SnFg 89 17. Mossbauer Data for SnF^~ Compounds 93
18. AsFo and SbF^ Complexes of Fluoro- and Oxy-element Heterocations 97
19. The Vibrational Spectra of ClF2SbF6, ClF2AsF6 and BrF„SbF, 100 26 + +
20. Vibrational Frequencies of the species C1F2 and BrF2 106 21. Frequency Assignments of AsF^ in ClF^sF^ 108 22. Fundamental Frequencies for the AsF^ Anion 109 23. Correlation diagram for the >*XEg Species 113 24. Fundamental frequencies for the SbF^ anion 114
25. Vibrational frequencies for the (Hal(S03F)^) Anion 125
26. Vibrational frequencies for Br(S03F)3 and I(S03F)3 126 vii!
Page
27. Correlation Diagram for the SO^F group 130
28. Vibrational Spectra of I0F3 and KIO^ 143
29. Vibrational Spectra of I02F, I02AsF6 and 1^ 147 30. Vibrational Frequencies of Bridging SO^F Groups 154
31. Specific Conductivities of I02S03F in HS03F 157
32. Specific Conductivity of Neat SeO(S03F)2 as a Function of Temperature 166
33. Specific Conductivities of SeOCl2 and SeO(S03F)2 in HS03F 171
34. Conductimetric Titration of SeO(S03F)2 in HS03F with Se 175
35. Specific Conductivities of SeO(S03F)2 in Superacid 178
36. The Raman Vibrational Spectra of SeO(S03F)2 and
SeO(S03F)Cl 182
37. Raman Spectra of the Solutions: SeOCl2 inlHSC^F, and
SeOCl2 in SeO(S03F)2 186 38. Chemical Shifts of some Oxy Selenium-Compounds 187 39. Se-0 and S-0 Stretching Force Constants 189
LIST OF FIGURES
1. Reaction Vessel for Solid-Liquid and Liquid-Liquid Reactions 12 2. Monel Metal 2-part Reaction Vessel (Front View) 13 3. Apparatus for the Preparation of $2®6^2 16 4. Electrical Conductivity Cell 22 5. 3 - Compartment Diaphragm Migration Cell (Front View) 36
6. Conductivities of Fluorosulphates in HS03F 48
7. Conductivity Curve of (CIO )S3010 in H2S04 50
1 8. H NMR Chemical Shifts for Solutions of G,102S03F in HS03F 52
9. Uv-visible Spectrum of C102S03F in HS03F 54 viii
10. The Walsh Diagram for Triatomic Species 56
11. Infrared Spectrum of C102AsF6 65
12. The Raman Spectrum of C102AsF6 66 13. Conductivity Curves of AsF^ Compounds 76 14. The High Resolution Infrared Spectrum of I^SnF^ 85
15. The Infrared Spectrum of (C102)2SnF6 87
16. The Raman Spectrum of (C102)2SnF6 88 17. The Mossbauer Spectrum of (ClO^SnFg 94
18. The Infrared Spectrum of ClF2AsF6 101 19. The Raman Spectrum of ClF^AsF,, 102 ZD
20. The Infrared Spectrum of ClF^SbF, and BrFnSbF^ 103 ZD ZD
21. The Raman Spectra of ClF„SbF^ BrF„SbF. and C10oSbF^ 104 ZD ZD ZD
22. The Crystal Structure of BrF2SbF^ 111
23. The Raman Spectra of KBr(S03F)4, KI(S03F)4 and Br(S03F)3 128
24. Apparatus Used for the Preparation of I0F3 139
25. Raman Spectra of I0F3 and KI02F2 144
26. Raman Spectrum of I02F 148
27. Raman Spectrum of I02S03F 149
28. Infrared Spectra of I0oF and I0„AsF, 150 f- Z ZD
29. The Proposed Structure for I02S03F 156
30. Conductivity Curve of I02S03F in HS03F 158
31. Conductivity of Neat Se0(S03F)2 as a Function of Temperature 167
32. Conductivity Curves of Se0Cl2 and SeO(S03F)2 in HS03F 173
33. Conductimetric Titration of Se0(S03F)2 in HS03F with Se 176
34. Conductivity Study of Se0(S03F)2 in HSbF2(S03F)4 179
35. Raman Spectrum of SeO(S03F)2 180 36. Raman Spectrum of Se0(S0 F)C1 181 ix
ACKNOWLEDGEMENTS
I am very grateful to my research director, Dr. F. Aubke, ;'7for his guidance and teaching during my years of graduate studies, and for his tremendous assistance in this research project.
I am also grateful to my colleagues with whom it has been a pleasure to work, and, in particular, Mr. A.M. Qureshi,;for many interesting and helpful discussions. I am indebted to Dr. J.R. Sams and his research group for obtaining the Mossbauer spectra, and to Dr. J.A. Ripmeester for 77 the Se NMR spectra.
Thanks are due to the staff of the machine shop, electronics department and glassblowing shop for the construction of much of the apparatus used in.this research, and to Mrs. B. Krizsan for the drawing of the spectral figures.
Finally, I would like to thank my parents for their encouragement, patience and understanding over all the years. 1
CHAPTER ONE
INTRODUCTION
1.1 OXYFLUORIDES OF THE HALOGENS AND SELENIUM
Element-oxygen-fluorine compounds of the type EOnFm where n and m are found to range from 1-5, and where both fluorine and oxygen are directly bonded to the element are commonly known as oxyfluorides.
Their physical and chemical properties depend primarily on the electro• negativity and the oxidation state of the element serving as the central atom.
The known oxyfluorides of the halogens and selenium are listed in Table 1 along with some of their properties and preparative routes.
The most general synthetic routes to oxyfluorides are:
(a) Fluorination of an oxide, e.g. CIO2F can be prepared by the fluorination of CIC^ •
(b) Fluorination of oxyanions. e.g. CIO2F can be prepared from
2 3 the fluorination of KC103 by BrF3„;or C1F3 ' . Also, fluorosulphuric acid has been found to be particularly useful in producing oxyfluorides
4 through its reactions with oxyacids and their salts .
(c) Partial replacement of fluorine in fluorides by oxygen, e.g. The 5-7 partial hydrolysis of IFy results in the production of IOF5 (d) Ligand redistribution between a fluoride and an oxide of a common 8-10 element, e.g. The reaction between 12^5 and IF5 produces IOF3 Review articles describing the preparations and properties of 11 12 oxyfluorides of the halogens and groups VB and VIB elements are available ' 2
TABLE 1 - OXYFLUORIDES OF THE HALOGENS AND SELENIUM
OXYFLUORIDE PREPARATIVE ROUTES PHYS. PROP.* REF. REMARKS
-115° 1 Highly reactive; C102F a. CIO 2i + F2 (-50°) m.p. b. KCIOT + BrF. b.p. - 6° 2,3 Attacks glass
+ F C103F a. KC103 2 (-40°) m.p. •146° 13-16,2 Relatively unreactive; b. KCIO4 + HSO3F (70°) b.p. • 47° 17 Very stable
Br02F a. Br02 + F2 (-50°) m.p. - 9° Thermally unstable; dec. 56° Attacks glass b. KBr03 + BrF5
I02F Decomposition of IOF3 at 110° dec. 210c 8-10 Stable in the absence of moisture
IO3F HIO4 + F2 (HF) dec. 100° 19 Stable in glass vessels.
I0F: I205 + IF5 dec. 110° -10 Crystal structure known
t.p. 4.6° I0Fc Partial hydrolysis 5-7 Very reactive of IF7 subl.-1.9°
c H0I0F4 + S03 m.p. 41 20 Stable in dry air; I02F3 2 different isomers
c SeOF. Se02 + F2 m.p. 15 30 Colourless liquid b.p. 126c 21 b. Se02 + SeF4
Se02F. c BaSe03 + HS03F (150°) m.p. -99 22 Colourless gas b.p. - 8C
All temperatures in degrees Centigrade 3
This present research is concerned with the potential of oxyfluorides to form oxycations and oxyfluorocations. This can be accomplished in two ways:
(a) The fluorine atom is abstracted from the oxyfluoride by a Lewis acid. This reaction occurs in general for a number of the oxyfluorides of the main group elements. For example, the reaction of AsF^ with oxyfluorides of nitrogen proceeds as follows:
+ FNO + AsF5 = N0 AsF6" (23)
+ FN02 + AsF5 = N02 AsF6" (24)
+ F3NO + AsF5 = F2NO AsF6" (25,26)
In particular, the formation of oxycations of the halogens and selenium might be achieved by the reaction of oxyfluorides with Lewis acids such as
AsF^, SbF^, BF3, S03 and SnF,^. (e.g. The reactions of these Lewis acids
with C102F leads to the complexes C102AsFQ, C102SbF6, C102BF4,
+ C102S03F and (C102)2SnFQ, but the existence of C102 has not been determined.)
(b) The oxyfluorides are solvolyzed in strong protonic solvents such
+ as HF, H2S04 and HSO3F to yield oxycations. (e.g. The existence of I02
in liquid HF with I02F as the solute has been proposed by Schmeisser 29 and Lang .)
A combination of both methods of forming oxycations would be the solution;-studies of oxycation complexes. Here, it would be advantageous to employ element oxyfluorosulphates as solutes and fluorosulphuric acid,
HSO3F, as the solvent for the following reasons: 4
(a) HSO^F is a strong protonic acid, and many polycations such as lg"1" and Se^** have been found to exist in this medium^. The acidity of the acid may also be increased by the addition of SO3 and SbF,. resulting 28 in superacid media
(b) Solution studies are simplified when the solute and solvent have common anions. (c) Many methods are available for the preparation of fluorosulphates.
The intermediate formation of the fluoride is not required (see section 1.3).
1.2 SCOPE OF THE PRESENT RESEARCH
The problems in this research project can now be defined as follows:
(1) the synthesis of oxyfluorosulphates of the halogens and selenium, analogous to the known oxyfluorides, and an examination of their physical and chemical properties. (2)". the identification of oxycations possibly existing in these compounds, and also in polyfluoroanion complexes of the Lewis acids ASF5, SbF$ and SnF^ either in solution or in the solid state. (3) spectroscopic studies of the related compounds in (2) containing fluorocations.
Though many of these compounds have been previously synthesized, structural studies for the majority are not available. 5
1.3 PREPARATIVE ROUTES TO FLUOROSULPHATES
Fluorosulphates were originally prepared by the addition of sulphur trioxide to element fluorides as according to:
EFn + nS03 » E(S03F)n
This reaction is useful in the preparation of alkali or alkaline earth metal fluorosulphates, but, in general, is not a suitable method for the synthesis of fluorosulphates. In the case of transition metal 32-34 and non-metal fluorides , for example, the addition of sulphur trioxide is often incomplete even at high temperatures and withjlarge excess* of SO^. In recent years, however, the syntheses of peroxydisiilphuryl difluoride (by Dudley and Cady^ in 1957) and its derivatives have made possible the preparation of a large number of fluorosulphates.
To date, there are three fluorosulphonating agents which are commonly employed to prepare fluorosulphates. The most effective agent
and most convenient to handle is peroxydisulphuryl difluoride, S20oF2-
Prepared by the catalytic fluorination of sulphur trioxide vapour at
180° in the presence of silver CH) fluoride, it is the fluorosulphato analogue to elemental fluorine. Indeed, its reactions are very similar 36 37 to those of fluorine where free radicals are involved :
2F- AH, = 38 kcal/mole Diss.
AH, = 21 kcal/mole- UlSS . 6
Peroxydisulphuryl difluoride behaves as a pseudohalogen whose reactivity and electronegativity is considerably greater than chlorine but less than fluorine. $20^2 displaces oxygen and chlorine from oxides and chlorides respectively to form the corresponding fluorosulphates, and adds across the double bond of perfluoro-olefins . It reacts 41 42 with the halogens to form halogen fluorosulphates ' which are similar in a formal sense to the interhalogens.
As in the case of fluorine, 320^2 reacts explosively with all organic matter igniting it spontaneously. As a consequence of
its physical properties (b.p. 67.1°, m.p. -55.4°), however, S20^F2 can be conveniently handled in a vacuum line, and in the absence of moisture, can be stored in glass vessels.
Although somewhat less easier to handle in a vacuum system, bromine monofluorosulphate ^>^, BrSO^F (b.p. 117.3°), has also proved to be a good fluorosulphonating agent. It is particularly useful in its reactions with non-metal(C,S,Si,P) oxychlorides and chlorides to yield bromine monochloride and fluorosulphates. By the use of bromine fluorosulphate as the fluorosulphating agent, oxidation of the non-metal, often encountered when S20^F2 is employed, is avoided.
The third fluorosulphating agent which has found use in the past in the preparation of fluorosulphates 6 The above discussion has meant to serve as an introduction to the methods of preparing fluorosulphates and the three fluorosulphating agents which are commonly used. Review articles discussing the chemistry of fluorosulphates^, and sulphur-fluorine compounds can be found in the literature^. The investigation of possible fluorosulphato compounds has been pursued in the past on the basis of the well-established chemistry of the analogous inorganic fluorides (e.g. S20^F2 v.s. F2, HSO3F v.s. HF and KBr(S03F)4 v.s. KBrF^). As synthetic routes to oxyfluorides (Table 1) often involve the use of elemental fluorine, it was hoped to prepare oxyfluorosulphates by analogous reactions substituting 820^2 for F2. 1.4 PREPARATIVE ROUTES TO POLYFLUOROANION COMPLEXES Polyfluoroanion complexes can most often be prepared by the reaction of an element fluoride or element oxyfluoride with a Lewis acid such as AsFg, SbFc; or BF3. This direct reaction of the fluorides is not always, however, the best means of synthesizing heterocation polyfluoroanion complexes. Other methods are available, and these can be summarized as follows: £a) The reaction of an oxyfluoride with the chlorine analogue of the Lewis acid. e.g. The synthesis of (ClO^^^F^ is achieved by the reaction of CIO2F with SnCl4- (b) The reaction of an oxychloride with the Lewis acid. e.g. The syntheses of N0AsF6, NOSbFg, N02AsF6 and N02SbF6 can be accomplished by reactions of N0C1 and N02C1 respectively with the 53 Lewis acids AsF,- or SbF5 . The highly reactive compounds NOF and NO2F are not required. (c) The reaction of an oxyfluoride with oxides. e.g. The reaction of C102F with Sb205 or at -10° leads to the 54 formation of C102SbF6 or C102BF4 . It is noteworthy that these three alternative methods of preparing polyfluoroanion complexes are essentially the same as the direct reaction of the fluorides since the oxides, chlorides and oxychlorides involved are fluorinated in the initial stages of the reaction. 9 CHAPTER TWO GENERAL EXPERIMENTAL TECHNIQUES The purpose of this chapter is to discuss general experimental techniques, apparatus, instruments and materials commonly used throughout the research described in this thesis. More specialized techniques and all chemical reactions other than the synthesis of ^2^6^2 are discussed separately in the appropriate chapter. 2.1 VACUUM SYSTEMS AND APPARATUS 2.1.1 General Nearly all compounds encountered in this research were unstable to moist air. It was necessary, therefore, to use a dry-box containing a purified nitrogen atmosphere for the transfer of materials. All volatile compounds were handled in either glass or monel metal vacuum lines constructed on metal frame-works inside fumehoods. Many of the compounds were found to be explosive towards all organic material, and some,particularly the selenium compounds, ' iare highly toxic. Consequently, great care was exercised in their handling, and all reactions were performed in we11-ventilated fumehoods. 2.1.2 The Dry-box The Vacuum Atmospheres Coorporation Model HE-43-2 Dri-Lab with a Model HE 93-B Dri-Train was used. Purified "L grade" nitrogen provided an inert atmosphere for the dry-box. 10 The nitrogen was circulated through the dri-train equipped with Linde's molecular sieves, which could be regenerated by use of an oven incorporated in the system. The process of regeneration was performed monthly to maintain a dry atmosphere. 2.1.3 The General Purpose Vacuum Line This consisted of a glass manifold with a B19 joint attachment to the safety trap assembly and five B10 outer joint outlets closed to the line by Fischer and Porter teflon valves. These greaseless valves were found to maintain a good vacuum, and in this respect, were superior to normal glass valves which often leak as a result of reagent attack of the stopcock grease. Vacuum was obtained by a standard rotary vacuum pump. The pressure of the system was measured by a mercury manometer attached to one of the B10 outer joints. A layer of Kel-F oil covered the mercury column to protect it from corrosive material. Before use, the vacuum manifold was cleaned with distilled water, dried, and finally flamed with an oxygen-gas burner to eliminate all traces of moisture. 2.1.4 The Metal Vacuum Line In the handling of compounds reactive towards glass, it was necessary to use a monel metal vacuum line. The vacuum manifold consisted of monel metal tubing, 1/4" O.D. and 1/32" wall thickness, joined by Swagelok fittings and equipped with Whitey valves (1KS4.316) leading to the outlets. 11 The manifold was attached to the glass safety trap via a "Cajon" joint. Gas pressures within the system were measured with a helicoid test guage and a Kontes ionization guage. 2.1.5 Glass Reaction Vessels A typical reaction vessel employed in reactions involving a non-volatile substance and a volatile reactant is shown in Figure 1. The starting material was introduced in the dry box at the B19 joint. After removal from the dry-box, the vessel was sealed at the constriction. The reaction vessel was then attached to the vacuum line at the BIO joint and the volatile reactant condensed on to the starting material by vacuum distillation. 2.1.6 Metal Reaction Vessels Compounds reactive to glass were handled in monel metal 2-part reaction vessels (Figure 2). These metal containers were equipped with a Hoke valve (No. 431), a removable lid and a teflon gasket which was fitted in between the lid and the vessel to provide a vacuum-tight seal. The lid was secured to the bottom vessel by six bolts fastened by an Allen wrench. The reaction vessel could be connected to one of the vacuum line socket outlets by means of Swagelok nuts, with front and back teflon ferrules inserted to provide a vacuum-tight fit. In addition, 1 litre and 500 ml monel and nickel containers, equipped with Hoke 431 or Whitey valves were used as storage or weighing containers for volatile fluorine compounds. B19 Cone + +- BIO Cone <- BIO Cone •«- Fischer and Porter Teflon Valve Teflon Thick-wall Glass Valve 125 ml Pyrex Erlenmeyer -> Flask Reaction Vessel Side View of Fischer and Porter Teflon Valve FIGURE 1 - Reaction Vessel For Solid-Liquid and Liquid-Liquid Reactions Hoke Valve (No. 431) Monel Metal Tube Bolts to Secure Lid to Bottom Vessel Condenser Inlet Bottom -«- Condenser Inlet Monel Metal Reaction Vessel (150 ml) FIGURE 2 - Monel Metal 2-Part Reaction Vessel (Front View) 14 2.1.7 The Fluorosulphuric Acid Distillation Line The apparatus and techniques involved in the purification of fluorosulphuric acid suitable for conductivity measurements have beeu n previousli y Adescribed -u J . 45,60 Fluorosulphuric acid was purified by double distillation, and for its use in conductivity measurements, finally distilled directly into the conductivity cell attached to the distillation apparatus'^ by means of a B19 joint. L-grade nitrogen was allowed to pass through the conductivity cell and distillation apparatus for at least two hours, during which time the apparatus was flamed periodically to eliminate moisture. Purified fluorosulphuric acid possessing -4 -1 -1 a specific conductivity less than 1.30 x 10 0, cm was deemed satisfactory for conductivity measurements. The lowest specific conductivity reported is 1.08 x 10 4 £2 •''cm * 2.1.8 The Fluorine Line And The Preparation Of S^F.^ The fluorine line and the flow apparatus used for the synthesis of ^2^(^2 (illustrated i-n Figure 3) was built from copper tubing (1/4" O.D., 1/32" wall thickness) on a metal framework in a fumehood. All connections were made by Swagelok fittings. Autoclave engineering valves were used to regulate the fluorine pressure from the cylinder. (A Crosby pressure guage indicated this pressure.) As shown in Figure 3, Hoke diaphragm valves (No. 413K) were used extensively throughout the system and were found to be reliable and resistant towards fluorine. Fluorine was passed through a stainless steel cylinder containing sodium fluoride to remove impurities of HF. Strong bolts which could be opened by an Allen wrench fastened the two removable lids to the cylinder, and leak-tight seals were achieved by teflon gaskets. The cylinder was wound with Nichrome ribbon, insulated with asbestos paper and finally with asbestos string to permit electrical heating, needed for the regeneration of sodium fluoride. The sulphur trioxide, contained in a pyrex flask, was carried by dry L grade nitrogen to the reactor where it mixed with a stream of fluorine. The reactants passed through the furnace forming products which were condensed out in traps A, B and C. A fluorolube oil bubble counter gave a visual indication of the flow rate of excess T^- The synthesis of ^2^(^2 ^aS ^een Previ°usly described in detail^ While the design of the apparatus is essentially the same as that reported in the literature, a number of modifications were made with the result of increased yields. The catalyst AgF2 was prepared in situ as according to Shreeve and Cady^"*". However, a larger reactor (120 x 7.6 cm v.s. 90 x 7.6 cm. was employed to extend the residence time of the reactants. Also, a reactor temperature of 180°C. (v.s. 150°C.) was found to give the best results. The reactor was heated in a similar fashion as the NaF trap, but it was necessary to wind a second coil of nichrome wire about the reactor to provide adequate heating. Thermocouples inserted beneath the asbestos layers indicated the1 temperature at various points along the reactor. Crosby Pressure Guage To cylinder To Soda-lime Trap Fluorolube Oil Tube Whitey Valve Hoke 413 Valve = Autoclave Engineering Valves FIGURE 3 - Apparatus for the Preparation of S n,F 2 6 2 17 To avoid the condensation of fluorine fluorosulphate (explosive in large quantities), traps B and C were both cooled at -78°C. (dry ice) instead of liquid oxygen (-183°C.) for trap C used in the literature preparation. Trap A was left empty and served to remove non-volatile side-products. Finally, sulphur trioxide was heated in its flask at about 50°C. (v.s. 25°C). A pressure of 15 mm in excess of atmospheric pressure for both F^ and resulted in the production of ^>2^(^2 at ab°u1: 150 g./hr. (v.s. 6 g./hr.) once the above reaction conditions had been obtained. A number of difficulties in this preparation may be avoided if the following points are noted: (1) The nitrogen inlet tube to the sulphur trioxide flask container should not be allowed to dip into the liquid, as liquid (y-SO^) tends to change to solid (g-SO^) which can result in the tube becoming plugged. (2) The copper tubing connecting the outlet from the reactor to the cooling traps must be flamed regularly to avoid clogging by condensed material in the tube. (3) Traps B and C must be checked regularly and replaced by empty traps when the $2®^2 ^eve^ rises t0 the inlet tube. (4) Also, large amounts of solid material (SO^) in these traps indicates that the flow rate is too high, while the presence of a green colour in the liquid (due to FSO^F) implies that the F2flow rate must be reduced. Unreacted sulphur trioxide was removed from the product by extraction with 96-98 % H»S0. in a separatory funnel in the fumehood. 18 (This operation should be only attempted when necessary. A well-ventilated fumehood is essential. Even though no reports on the toxidity of 6 2 S^O^F^ exist, the reported extreme toxidity of the related S^O^F^, disulphuryl difluoride, can be regarded as a warning.) The peroxide fraction was then cooled to -78°C. and pumped for several hours until all traces of fluorine fluorosulphate had been eliminated. Finally, S^O^F^ was vacuum distilled into a teflon valve storage trap. Its purity was checked by an'IR spectrum of the gas and an NMR spectrum of the liquid. 2.1.9 Source of Materials and Equipment This section gives further information and the names and addresses of manufacturers of apparatus, equipment and materials used in this research: 1. Kel-F grease: Formula = Cl (C^-CFCl^Cl; 3M, Minnesota, Mining and Manufacturing Co. St. Paul, Minn. 2. Fluorolube grease GR-362: Fisher Scientific Co., Fair Lawn, N.J. These greases were used successfully in the lubrication of stop-cocks and apparatus employed in the handling of fluorine-containing compounds. 3. The Crosby High Pressure Guage: Manufactured by Crosby, Ashton Valve and Guage Co. , Wrentham, Mass. 4. Kontes Ionization Guage: Kontes of Illinois, Franklin Park, -3 Illinois 60131. Useful in measuring pressures in the range 1mm -10 Hg. 5. Helicoid Test Guage: American Chain and Cable Co. Bridgeport, New Jersey: Useful in measuring pressures in the range 1mm - 1000 mm. 19 6. Swagelok Fittings: distributed by Crawford fittings (Canada) Ltd. Niagara Falls, Ont. 7. Fluorolube Oil: formula.= CF2C1 (CF2CFCl)xCF2Cl; Hooker Chem. Corp. Niagara Falls 8. Linde Molecular Sieves: produced by Linde Air Products and distributed by Fisher Scientific Co. 9. Fischer and Porter teflon valves: Fischer and Porter Co. Warminster, Pa. 10. Kontes Valves: Kontes of Illinois, Franklin Park, Illinois 60131 11. Rotaflow valves: Made by Quickfit and Quartz Ltd., Stone, Staffordshirey England; Valves in 9, 10, and 11 incorporated in glass vessels were found convenient to use and leak-tight. 12. Whitey valves (1KS4.316): manufactured by Whitey Research Tool Co. 5525 Marshall St., Oakland 8, Calif. 13. Hoke Valves: (a) No. .431 (b) No. 413K Made by Hoke Inc. Creskill, New Jersey 14. Autoclave Engineering Valves (MV30): made by Autoclave Engineering Inc. 2930 W. 22nd St.,Erie, Pa. 20 2.2 REAGENTS A number of purchased reagents which were used extensively during the course of this research are indicated in the table below. Other chemicals used less frequently and not listed here, and general "shelf" reagents are described in the experimental section under the appropriate chapter. COMPANY SUBSTANCE % PURITY CLAIMED BY MANUFACTURER Allied Chemical S03 (Sulfan) Not Given (Canada) Ltd. Not Given HS03F Not Given 95.5 - 96.5 H2S04 Alpha Inorganic SbF5 Not Given Chemical Corp. KAsF^ 6 Not Given KSbF5 95 K2SnF6 Not Given British Drug Houses Se 99.0 Se02 99.0 SeOCl. 97.5 Matheson Co. HF 99.9 C1F, 98.0 BrF_ 98.0 Ozark Mahoning Co. Not Given 2.3 PHYSICAL EXPERIMENTAL METHODS 2.3.1 Electrical Conductimetry Solutions of a number of compounds in fluorosulphuric acid and sulphuric acid were studied by electrical conductivity measurements. 45 The techniques involved, including the design of the conductivity cell , 45 the purification of HSO^F , and the preparation of "minimum conducting" H^SO^ '^^, have been previously described. A three-electrode conductivity cell with approximate cell constants of 5, 10 and 15 cm * was used. The cell was calibrated with aqueous solutions of potassium chloride^, and the electrodes' were platinized^, as routine procedures preceding a conductivity run. Conductivity measurementsjwere made witL a Wayne Kerr Universal Bridge B 221A, while the cell was contained in a constant temperature oil bath of 25.00'± 0.01°C. with a Sargent Thermonitor CModel ST). The bath temperature was measured with a precision thermometer, previously checked against a platinum resistance thermometer, obtained from Brookly Thermometer Co. The conductivities of neat liquids were measured in the electrical conductivity cell illustrated in Figure 4. The cell constant was about 1 cm-*, and approximately 2 ml. of liquid were required for accurate conductivity measurements. The cell, loaded in the dry-box, contained a purified nitrogen atmosphere. 22 FIGURE 4 - Electrical Conductivity Cell (actual scale) B19 Copper Wire 1 f Mercury Sample pt Solid Glass Electrodes 23 2.3.2 Infrared Spectroscopy Infrared spectra were recorded on a Perkin-Elmer 457 Grating Spectrophotometer covering the range 4000 - 250 cm ^. Higher resolution infrared spectra were obtained from a Perkin Elmer 301 High Resolution Grating Spectrophotometer (4000 - 100 cm *). Owing to the high reactivity of the majority of compounds investigated, either silver chloride or KRS-5 windows were used without the use of mulling agents. KRS-5 (Thallium bromide-iodide) windows have the advantage of permitting the recording of infrared spectra down to 250 cm 1, but are generally more susceptible to chemical attack than are the AgCl windows which "cut off" at about 400 cm Gaseous infrared spectra were obtained using a monel metal cell equipped with a Hoke valve (No. 431) and fitted with vacuum tight AgCl windows. The cell windows were cut from AgCl sheets (0.042" in thickness) and placed in the cell in the absence of light. Black tape was employed to protect the windows when the cell was not in use. All infrared optical windows were obtained from The Harshaw Chemical Company, 1945 East 97th St., Cleveland, Ohio 44106 2.3.3 Raman Spectroscopy A Cary 81 Spectrophotometer equipped with a Spectra Physics Model 125 He - Ne laser as a source of exciting light (A. = 6328 A) was used. The solid samples were contained in "optical" flat bottom pyrex tubes (6 mm O.D.). 24 Raman cells for liquids were similar to the cells used for solids but were either T-shaped orL-shaped to permit their being flame-sealed after being loaded and capped inside the dry box. 2.3.4 Mossbauer Spectra The Mossbauer spectrometer consisted of a TMC Model 305 velocity transducer driven at constant acceleration by a TMC model 306 wave form generator which also synchronized the 400 channel analyser. The transmitted radiation was detected by a Reuter-Stokes RSG - 60 proportional counter with 2 atmospheres Xe-CH^ as fill gas, and fed into a Nuclear Chicago model 33-15 single channel analyser and then to a 400 channel memory. The output was displayed on an Hewlett Packard model 120 B oscilloscope and also printed out on a model 44-16 IBM typewriter. The printout of the spectrometer was fitted to a Lorenztian curve on an IBM 4070 computer. Other computer programmes were available to possibly resolve quadrupole splitting in broadened peaks. The source of the 119m gamma radiation was BaSnO^ enriched with Sn which provides a single narrow emission line. Samples were placed on a brass cell with mylar windows and the spectra run at 295° and 80°K. The doppler velocity scale was calibrated against the quadrupole splitting of an NBS standard nitroprusside absorber. The isomer shifts were reported relative to Sn02 at 80°K. The estimated precission of Mossbauer parameters is ± 0.03 mm sec l. 25 2.5.5 UV - Visible Spectroscopy A Cary Recording Spectrophotometer Model 14 with a lamp supply (IR, UV, Visible) Model D from Applied Physics Corporation was used to obtain electronic spectra in the range X = 2000 - 7000 A. Quartz optical cells were obtained from Pyrocell Manufacturing Co. and provided with optical spacers which reduced the path length from 10.00 mm to 1.00, 0.50, 0.20 and 0.05 mm. A teflon gasket fitted between the spacer and the cell excluded moisture from the sample solution. The cells were kept tightly shut to prevent fogging,and hydrolysis during the measurement by using a cell holder made in our machine shops. The solutions were prepared and the cells filled in the dry box. 2.3.6 Nuclear Magnetic Resonance Spectroscopy 19 F NMR spectra were obtained,using a Varian HR100 spectrometer, and employing CFCl^ in glass capillaries as an external standard. ^H NMR spectra were measured with respect to HSO^F as an external standard, and by using a Varian HR60 spectrometer. 77 Se NMR spectra were recorded on a Varian DP60 spectrometer. A description of this instrument and the techniques used in obtaining 6 7 the spectra have been previously described with the following exceptions: 77 The Se resonances were observed at a frequency of 10.0 MHz in a resonance field of about 12300 gauss. For narrow lines, the method of Acrivos was used with a modulation frequency of 432 cps and a modulation index of 1.5 - 2.0, while for the broad lines, the derivative method was employed. Samples were placed in 15 mm O.D. tubes. Chemical shifts were obtained by the sample replacement method, using neat SeOC^ as a reference. 2.3.7 X-Ray Powder Photography Finely powderedvsamples were packed into 0.5 mm Lindemann glass capillaries inside the dry-box. The tubes were sealed with Kel-F grease and then flame sealed after removal from the dry-box. X-ray powder photographs were obtained using a General Electric Powder Camera of 14.32 cm diameter and having the conventional Straumanis arrangement. Time of exposure using nickel-filtered copper x-radiation (XK^ = 1.54050 A) generally varied from 4-6 hours. Illford "Illfex" film was used for the spectra. A film illuminator and measuring device obtained from Philips Electronics Inc. was used to measure the line spacings from which "d" spacings could be obtained. CHAPTER THREE CHLORYL FLUOROSULPHATE CHLORYL COMPOUNDS (RED) PART I 3.1 INTRODUCTION Chloryl fluorosulphate, C102S0gF, was first prepared in small 3 amounts by Woolf in 1954 by the direct addition of sulphur trioxide to chloryl fluoride. This reaction was described to be violent, and subsequent investigations on the reactions and properties of chloryl fluorosulphate were not made. Woolf postulated an ionic structure for this compound, 002-SOgF . Chloryl fluorosulphate is a highly reactive red solid at room temperature which melts at about 27°C. However, it is often observed as a red liquid at room temperature due to the presence of small impurities, and a tendency to supercool, common to many fluorosulphat derivatives. A related compound and also a red solid, dichloryl trisulphate, (0102)2^3020? was prepared by Lehmann and Krueger by the reaction of SO^ and C^Oj."'''. By analogy to the ionic compound, (N02+)2^3^10^ ^ dichloryl trisulphate was also suggested to have an ionic structure. No structural investigations have been made, however, to support the ionic formulations for either CIO2SO2F or (C102)2^30^^. 28 It is the object of the research described in this chapter to find convenient routes to the preparation of chloryl fluorosulphate, and to prove the existence of the chloryl cation, VAO^. 3.2 EXPERIMENTAL 3.2.1 The Reaction Between KC1C>3 and S206F2 An excess of S20^F2 was distilled in vacuo to a reaction vessel (Figure 1) containing, in a typical reaction, 31.52 mmoles of finely powdered KCIO^ (A.R.) at -196°C. Upon warming to room temperature, a mild reaction took place with the liberation of 02, and the formation of a red liquid and a white solid. (At the temperature of liquid nitrogen, the colour of the red compound faded to a yellow colour, a property characteristic of chloryl compounds.) The reaction rate was increased by heating at 50°C, and after about 4 hours, the evolution of 02 had ceased and the reaction was complete. After the removal of S205F2 by pumping, chloryl fluorosulphate was separated from the product mixture by distillation in vacuo and at elevated temperatures, about 80 - 100°C. The remaining product corresponded to 31.59 mmoles of KSO3F, and its identity was confirmed by an infrared spectrum. Chloryl fluorosulphate was found to be a very reactive liquid, igniting organic material and exploding with water. All attempts to prepare KCIO^SO^F^ at low temperatures by controlling the reaction of KCIO^ and $2^(3* 2 at a temPera1:ure °f -20°C. were unsuccessful. In all cases, a mixture of CIO2SO.JF and KSO^F resulted. The existence of the fluorine analogue, KCIO2F2, has 81 been recently reported by Fox et al. 29 3.2.2 The Reaction Between C102 and S206F2 Chlorine dioxide was prepared from KCIO^ and L^SO^ in the presence 57 of oxalic acid as according to Brauer . The purification of CIO2 was accomplished by pumping at -78°C. to eliminate impurities of C02 and by trap to trap distillation. As CIO2 is known to be an unpredictMely explosive compound, the following precautions were taken to minimize the dangers involved in its preparation and handling: (1) The preparation of CIO2 was performed in a darkened fumehood as CIO2 is light sensitive. (2) Only small quantitites of C1C>2 were prepared in a single run 57 (about 1/10 the amount of starting materials suggested by Brauer ). (3) The CIO2 generated in the reaction flask was immediately cooled by a cold water jacket incorporated in the stem of the reaction flask. An excess of S20^F2 was distilled on to about 3 g. of C102 contained in a reaction vessel cooled at -196°C. The reaction flask was warmed to -40°C, where a mild reaction took place resulting in the formation of red crystals of C102S03F. A check for the production of non-condensible gases was negative as the manometer registered zero pressure for the system at -196°C. The final product was obtained in a pure form by removing the excess peroxide by distillation at room temperature. Its existence as a red crystalline solid at room temperature - it melted when placed in its container in a water bath at about 30°C. - was indicative of the high purity of the product. 3® 3.2.3 Other Attempts To Prepare Oxyfluorosulphates Of Chlorine 58 An excess of 820^2, prepared by the reaction of SO3 with SbF^ , was distilled in vacuo on to KCIO^ contained in a reaction vessel. At temperatures up to 80°C, no reaction was evident. ^2^6^2 fa-i-'-ec* t0 react with ClO^F and KCIO^ up to temperatures of 100°C. 3.2.4 The Reaction Between NaC10„ and S„0.F„ I Z o I The reaction between finely powdered NaC102 CA.R.) and $20^2, performed under conditions similar to those in section 3.2.1, proceeded slowly at -10°C. and rapidly at room temperature with evolution of O2, CIO2 and the formation of a red solid. At reaction completion, the 19 excess peroxide was removed by distillation. Infrared and F NMR spectra revealed a large amount of ^>2^^2 to be present in the residual peroxide. The red solid melted at about room temperature, and its removal by distillation in vacuo left behind a white solid which was indicated by the weight change and an infrared spectrum to be NaS03F. In a different run, the reaction was halted at an incompletes stage and a Raman spectrum obtained of the red-coloured product mixture. The following products were identified: Unreacted NaC102 -1 1 (V3C1() - 832 cm , V1C10 - 804 cm" ), NaS03F (1291, 1078, 800, 592, 2 -1 2 -1 -1 570, and 420 cm ) and C102 (v3ci0 1103 cm , vlclQ 949 cm ). No evidence for the chlorosyl cation, CIO , isoelectronic with 0^ -1 59 + (vn = 1555 cm ) was found, while C10? vibrations were either hidden by SO2 vibrations, or did not appear as a result of absorption of the exciting light by the red coloured species. On one occasion, the reaction vessel exploded without warning during the course of the reaction. The explosive substance in the system was undoubtedly chlorine dioxide. Therefore, the above reaction should be performed with extreme care. 3.2.5 The Reaction Between KNO, And So0^F„ 3 2 6 2 The reactions of $2®^ 2 Wlth N0.j and NO^ were investigated to gain further information on the reactions of S^O^F^ with salts of oxyacids in general. An excess of S_0^.Fo was reacted with KN0_ (A.R.) under similar 2 6 2 o conditions as in section 3.2.1. The reaction proceeded rapidly at room temperature with the evolution of 0^ and the formation of a mixture of white solids identified by Raman spectra as KSO^F and N02S03F. The weight change was in agreement with the formation of these solids and 0^ as the only products. 3.2.6 The Reaction Between NaNO 'and S.0.Fo z 2 o z An excess of So0.F„ was reacted with NaN0„ (A.R.) under similar 2 6 2 2 conditions as in section 3.2.1. The reaction proceeded slowly with the evolution of O2 and formation of white solids. To complete the reaction, the flask was heated at 65°C. for several days. A Raman spectrum of the product mixture revealed the presence + 1 of NaS03F and N02S03F (N02 1408 cm" ). A second run, in which the S„0^F. was removed before reaction z o I completion and the temperature kept at 25°C, resulted in a different 1 Raman spectrum for the product mixture. Again, NaS03F (S02sym 1072 cm + was observed, but the peak at 1408 cm \ typical for N02 was absent. Instead, a peak was seen at 2303 cm 1, indicating the presence of N0+ -1 and therefore N0S03F (S02 sym 1090 cm ). Peaks observed at 1333, 1 832 and 1265 cm are due to unreacted NaN02- 3.2.7 Preparation of (ClO^^O and Dichloryl trisulphate and potassium trisulphate were prepared by the reaction of potassium chlorate with sulphur trioxide as according to Lehmann and Krueger."^ (C102)2S30^Q was separated from K2S30^Q by vacuum sublimation at temperatures less than 80° C. into a trap cooled at -78° C. To remove impurities of (C10)(C102)S3010, a small quantity of C102 was distilled on to the crude CC102)2S30^Q and the mixture allowed to warm to 0°C. over a period of several hours. The excess C102 and the liberated gases, 02 and Cl2 were removed by pumping at 0°C. The final product melted at 74° C. (reported 75.5°C.55). The residual K2S301Q was heated at 150°C. until all traces of (C102)2S30^Q had been eliminated. The resulting white solid was treated with an excess of S03 to convert impurities of IC^O^ , 33 arising from the decomposition of K^S^O^^, back to -^2^3^10* T^e ^3 was then removed by pumping, and the purity of the resulting product checked by an acid-base titration of the aqueous,-solution. 3.2.8 Reaction of CIO^SO F with KF At room temperature, chloryl fluorosulphate reacted vigorously with anhydrous KF, and with efforvescence. An infrared spectrum showed the gas to be ClO^F, while a Raman spectrum identified the residual solid as KSO^F. 3.2.9 Reactions of C102S03F with Perfluoro-Olefins 3.2.9.1 C2F Into a two-part metal reaction vessel (Figure 2), about 1 g. of ClO^SO^F was added in the dry-box. A small excess of C^F^ (Peninsular Chem. Research Inc.) was distilled on to the CIC^SO^F at -196° C. Upon warming to room temperature, a violent explosion erupted. No further studies of this reaction were attempted. 3.2.9.2 C.F. 4 4 An excess of perfluorocyclo-2-butene was distilled on to about 1 g. ClO^SO^F contained in a glass reaction vessel cooled at -196°C. Upon warming to room temperature, the reaction vessel and contents exploded. The study of the reaction between C^F^ and ClO^SO^F was subsequently abandoned. An excess of perfluoro-2-butene (Matheson Co.) was condensed into a heavy-walled 100 ml. reaction vessel, equipped with a Fischer and Porter teflon valve, containing about 2 g. C102S0gF at -196°C. At room temperature, two layers consisting of unreacted C102S0,jF (bottom red layer) and unreacted C.F (upper colourless layer) were 4 o observed. On warming to about 70°C. and shaking the container vigorously for about 20 minutes, all traces of C102S03F disappeared and a colourless liquid remained. The removal of the excess olefin left behind a viscous liquid. The weight change was considerably less than that expected for a 1:1 addition compound. Both infrared and Raman spectra of the liquid product showed the presence of covalent SO^F groups, and the absence of the olefin C=C double bond. No strong peaks in the region 850 - 1250 cm 1 were observed, and therefore, it was difficult to reconcile the existence of covalent CIO2 groups in the compound. In addition, a weak peak was seen in the vibrational spectra at 1890 cm 1 and it can be attributed to the FC=0 grouping. 19 Due to the complexity of the F NMR and mass spectra of the product mixture, the products were not identified. 3.2.9.4 CTF, o o Into a 500 ml glass reaction vessel containing about 1 g. C102S03F, 1 atmosphere of hexafluoropropene (Peninsular Chem. Research Co.) was condensed. An exothermic reaction took place at room temperature 35 between the gas and the red liquid. After the elapse of about % hour, all CK^SO^F had reacted, and a colourless liquid remained. All volatile material at -78°C. was distilled into another trap. Gaseous infrared spectra of the product were very similar to that obtained from the products of the C.F - CIO SO F reaction. Again, covalent S0_F groups 4 o z S 6 and the FC=0 group were seen, while the olefin C=C bond was absent. 3.2.9.5 C,F.._ o l U The reaction between ClO^SO^F and perfluorocyclohexene (Peninsular Chem. Research Inc.) was carried out in similar operations to those described in section 3.2.9.3. At room temperature, a slow reaction took place, but on heating for about 1 hour at 50°C. and with vigorous shaking, the red colour of ClO^SO^F disappeared and a colourless liquid remained. The excess olefin was removed by distillation at -10°C. Again, infrared spectra indicated compounds containing the FC=0 group, rather than 1:1 addition products. 3.2.10 Solution Studies 19 1 The behavior.of C102S03F in HS03F was examined by F and H NMR spectroscopy, UV-visible spectroscopy and conductivity measurements. In addition, conductivity measurements were obtained for (C102) 2S30.^Q and K2S301Q in HS03F and (C102)2S3010 in 100% H2S04. 36 The purification of HSO^F, referred to in section 2.1.7, required additional steps before its use in UV-visible spectral studies. About 2 g. S^O^F^ were added to the crude HSCLF, converting any impurities of SO (X = 2940 my) o z in 3.x to S„0oFo. Excess S„0,Fo and S„0oF_ were then eliminated in the forerun. 3o2 z o 2 o o z + Spectra of the C102 cation were obtained by running the sample spectrum versus neat HSO^F in the reference cell. A three compartment diaphragm migration cell used in studying (in HSO^F) the electrolysis of ClO^O^F is illustrated in Figure 5. The middle compartment was filled in the dry-box with a solution of ClO^O^F in HSO F and the outer compartments with HS0_F. After the current was turned on, the cathode compartment gradually turned yellow as chlorine dioxide was liberated at the cathode. The electrolysis was effected by a Fisher Scientific Electro Model D-612T filtered D.C. Power Supply. B19 Socket B19 Cone B19 B19 HS03F -* Sample Solution > Electrode Electrode Diaphragm FIGURE 5 - 3-Compartment Diaphragm Migration Cell (Front View) The addition of KCIO^ to HSO^F in a test tube produced a yellow solution with vigorous efforvescence of ClO^. It seems that chlorine dioxide is produced in the solvolysis of KCIO^ in HSO^F, rather than chloronium cations. This is in contrast to the solvolysis of nitrates and nitric acid which solvolyse in strong protonic acids to produce .„ •. 68-70 mtronium cations 3.3 RESULTS AND DISCUSSION 3.3.1 The Preparation of C102S03F; The Reactions of S206F2 With Chlorine and Nitrogen-Containing Oxy-salts Two convenient and novel routes to the synthesis of ClO^SO^F have been obtained: 25°C. KC10- + So0,Fo — • KS0_F + C10oS0_F + %0o j z o z Slow 3 2 3 2 -40°C. 2C102 + S206F2 2C10„S0,F Fast 2 3 Both reactions are analogous to the formation of C102F from the fluorination of KClOj2,3 and ClO^. Reaction (1) is typical of the chemistry of ^O^F^ where the preparation of fluorosulphates is accomplished by fluorosulphate free radical attack and subsequent expulsion of oxygen. The failure of S20;_F2 to react with KCIO^ is not surprising as a free radical mechanism is not available in its reactions. 38 The second reaction occurs at a faster rate and also eliminates the necessity of separating ClO^SO^F from solid KSO^F in (1). The main disadvantage here, of course, is the use of the potentially explosive CIO,,. The preparation of ClO^SO^F from SO^F- radicals and CIC^ was predicted by an earlier and analogous free radical 37 reactions by Lustig and Cady in which these authors prepared NF_S0„F from S.0,F_ and N0F.. Z O Z 0 Z Z 4 S„0,F„ failed to react with KC10. and C10-F, isoelectronic to the 2 o 2 4 J perchlorate anion. The stability of the coordinatively saturated •71 72 ClO^F molecule and the ClO^ anion is well-known. ' The reaction between NaCIO,, and S^O^F^ produces chlorine dioxide as a result of S^O^F2 behaving as a pseudohalogen: NaC10o + S_0,F- > NaSO-F + C10o Z Z D Z J Z C102 + %S206F2 • C1G2S03F C102 combines with fluorosulphate radicals to form CIC^SO^F. The liberation of CL may result from the oxidation of C1CL to C1CL by S OF., 2 l 6 2 '6 1 S F aS in A second which further reacts with 2°5 2 possibility as involves the replacement of an oxygen atom in C102 by SC^F' according to: CIO ~ + Sn0V^ • C10S0„F + S0„F~ z z o 2 JO Oxidation of either C10S03F or C102" by S206F2 to C102S03F and C103" respectively would account for the S,,03F2 found in the residual peroxide after the reaction. 39 However, no evidence was found to support the formation of ClOSO^F. The Raman spectrum failed to reveal any absorption bands in the frequency + -1 59 region expected for CIO , isoelectronic with 0^ (vo-0= cm ) • The overall reaction using an excess of ^20^^ appears to be: NaC10„ + 2So0,Fo(excess) • NaS0„F + ClO^SO^F + S.0_F_ + %0o / Z o Z j Z 6 Z o Z Z Peroxydisulphuryl difluoride reacted moderately with the analogous nitrogen-containing compounds, KNO^ and NaNO^ at room temperature as according to: KN0_ + So0.F9 • KS0„F + N0„S07F + %0. 5 Z o Z .5 Z o Z NaN0o(exc.) + So0£Fo NaSO F + N0S0_F + %0_ Z Z o Z >• J o Z NaN0„ + 2So0.F„(exc.) — • NaS0„F + NO.SO.F + S„0cF^ + h0„ z z o z o z o z D z z The reaction between $2®^ 2 and KNO^ is analogous to the KCIO^- F react n ^2^6 2 i° - However, NaN02 does not apparently react with $2®^ 2 in a similar fashion as its chlorine analogue, NaC102- Instead;; .^O^^ displaces an oxygen atom from the nitrite ion to form NOSO^F. In a controlled reaction, using an excess of NaN02, the Raman spectrum of the reaction product mixture reveals the presence of two fluorosulphates, + 1 + 1 NaSQ3F and N0S03F (N0 at 2303 cm" ); the N02 peak at 1408 cm" is absent. However, an excess of ^O^F,, in the reaction results in the production of N02S03F and NaS03F. As N0S03F is oxidized to I^SO,^ by the peak in the Raman spectrum at 2303 cm 1 diminishes, while the absorption peak at 1408 cm 1 due to the N0^+ cation becomes more intense. 40 Due to the difficulty in separating NOSO^F or r^SO^F from the alkali metal fluorosulphate, the above reactions are unsuitable preparatory routes to these fluorosulphates. Convenient syntheses for NO- and 73 74 NO2- fluorosulphates have been previously reported. ' 3.3.2 The Properties and Reactions of C102S03F At room temperature, C102S03F exists as a highly reactive supercooled liquid (m.p. 27°C.) , which explodes with cold water and organic material. It is thermally stable, however, up to temperatures of 150°C, and may be kept in glass storage vessels for long periods of time. In ionic solvents such as HSO^F, C102S03F is miscible, but it is only slightly soluble in the covalent liquids S^O^J^ and ^2^(^2' Electrical conductivity values for the neat liquid are given in Table 2: TABLE 2: Electrical Conductivity of C102S0 F TEMPERATURE (°C.) K (Q 1cm~1) 5.6 3.88 x loj 10.7 4.70 x 10 Jl 16.1 5.73 x loj 25.0 7.51 x 10 32.0 9.00 x 10 The relation of the specific conductivity to temperature is linear with a positive temperature coefficient, which is typical of an ionic liquid. The relatively high neat liquid conductivity (c.f. BrF^ 3 1 1 75 2 1 8.00 x 10~ fi~ cm" versus 7.51 x 10~ fi^cnf for C102S03F at 25°C.) + is attributed to self-dissociation: C10 S0„F = C10o + S0_F~. 2 3 2 3 Chloryl fluorosulphate appears to be a strong oxidizing agent. Its reaction with perfluoro-olefins result in oxidation and cleavage at the double bond in contrast with S„0,F„ which forms 1:1 addition z 6 z j 38 compounds Whereas ^2^(^2 ^a^*s t0 react with potassium fluoride, ClO^SO^F undergoes a vigorous reaction with efforvescence of ClO^F as according to KF + C102S03F •—• KS03F + C102F + Such a reaction suggests the potential use of C102S03F as a powerful fluorosulphating agent. + 3.3.3 The Existence of the Chloronium Cation, C102 , in Solution 3.3.3.1 Introduction T, . + £ . +. 76-80,82-85 M.+ The existence of the nitrosonium cation , NO , •j . .,. . 86-90,68-70 + . . . ' , and the nitronium cation ' in stronS protonic solvents and in the solid state, is well-known. These cations originate from the oxides NO and N02 respectively which are electronically similar to C102 in so far as they contain a single electron in the highest lying occupied molecular orbital (which is anti-bonding with respect to the central atom and oxygen.). A number of ionization potentials for simple diatomic and triatomic 48 49 molecules are listed in table 3 ' . While the first ionization potential of C102 is greater than that of both NO and NO,,, it is less 93 than that of 0„, which exists as the cation in the solid 0„PtF, Z ZD Thus, on the basis of ionization potential data, the existence of CIO* is quite possible . 42 TABLE 3: First Ionizationpotentials in ev •• a) diatomics NO NO + 9.35 o2 12.2 CO co+ 14. I Nt 15.7 b) triatomics N02 NO. 9.8 CIO. CIO + I I. I NO N20^ 12.9 2 SO. SO+ 13. I + CO, co2 13.8 43 The strongly protonic solvent, fluorosulphuric acid, has been 45 94 shown to be very useful as a medium for studying cationic species ' It undergoes autoprotolysis as according to: H S0 F+ + S0 F 2HS03F = 2 3 3 ~ Bases in this solvent are defined as substances which when dissolved in HSO^F increase the concentration of the fluorosulphate ion SO^F ; whereas, acids are defined as substances which increase the concentration of H^SO^F*. The solute KSO^F which behaves as a strong base in this solvent is often used as a standard in studying other solutes. The latter are compared by the calculation of y values, defined as the number of moles of fluorosulphate anions produced per mole of solute. As an approximation, y values are calculated by dividing the solute molality by the molality of KSUgF, the reference solute, at a given conductivity. 3.3.3.2 Conductivity Measurements Chloryl fluorosulphate dissolved readily in HSO^F to form a red coloured solution. Conductivity measurements were obtained for the dilute solutions, up to 0.045 molal ClO^O^F. The conductivity results, together with those obtained for (CIO2)2S30^Q, ^2^3^10 and the reference solute KSO^F, are listed in Table 4. As a means 29 of further comparison, the solutes NOSO^F and J^SO^F are included in the table. TABLE 4 (a): Specific Conductivities of CIC^SCLF, (C102)2S30 Q and K^O Q in HSCLF at 25.00°C. CIO tcio2)2s3o K S 2S03F 10 2 3°10 KSO-F Additions; 2 2 4 2 4 Molality x 102 K X 104 K X 104 Molality x io K X 10 Molality x 10 K x 10 Molality x 10 -1, (Q *cm ^) fn-l (fi cm j (fi ^cm (ft cm J 0 000 1.247 0 01032 95.69 0.0000 1.161 0.0000 1.236 0 060 1.778 0 1083 97.56 0.0098 1.456 0.0102 1.407 0 122 2.832 0 3973 102.5 0.0450 2.586 0.0317 2.158 0 161 4.036 0 9236 111.4 0.0921 4.746 0.3382 16.33 0 282 7.122 1 873 126.8 0.1640 8.234 0.4208 19.83 0 360 9.831 2 425 135.4 0.2131 10.62 0.7767 36.. 78 0 498 12.78 3 313 148.7 0.2707 13.48 0.8473 40.11 0 769 19.38 3 767 153.2 0.4721 22.92 1.0781 50.43 0 950 23.46 5 .279 172.8 0.5698 27.87 1.2851 59.91 1 241 29.85 0.7379>- y 35.72 1.5645 72.34 1 509 35.73 0.8249 39.92 1.8598 84.37 1 722 40.24 2 171 49.60 3 579 77.84 4 550 95.31 45 TABLE 4 (b): Interpolated Specific Conductivities of Solutes in HSO F (2S.00°C.) Molality KS03F (CIQ2) OSQ2F (N02) OSQ2F NO(OSQ2F) (CIO) s o 4 2 2 3 JO x IO2 K x IO K x IO4 K x IO4 K x ib4 K x IO4 1 X 1 X X Epr' cmi \pr' cm-'] [rr'cm- ] [p- cm~] [or cm'} X O.OO I..085 1.247 — 1 .116 — 1 .084 1 . 161 0.25 7. O 6. 46 0.91 4.87 0.70 6.1 0.87 12.45 1.78 O. 50 13. 61 12.8 O. 94 12.9 0.95 12.4 0.91 24. 3 1,79 O. 75 19.7 18. 9 0.96 19. O O. 97 18.8 0. 95 36.8 1.84 1 . OO 25.8 24.7 O. 95 25. 1 O. 97 24.6 0.95 48.4 1.88 1 . 50 38.0 35.5 0.93 37. 1 O. 98 36.8 O. 97 2.00 50.0 45.7 0.91 47.4 O. 95 50.9 1.02 2.50 62,5 57. 1 0.91 59.5 0.95 60.6 O. 97 3.00 72.7 67. 8 0.93 70. O O. 96 72. 2 0.99 3. 50 84.9 76.1 O. 90 81. 8 0.97 84.2 0.99 4.00 97. O 87.0 O. 90 90.0 O. 93 93. 3 O. 96 4.50 106. 8 94.3 O. 88 98.7 0.93 105. 2 O. 99 i i Chloryl fluorosulphate behaves as a strong base in HSO^F as is evident from the y values and a continuing increase in conductivity when KSO^F additions are made to the solution. The y values which are close to the reference KSO^F value of 1.00 indicate C102S03F to be completely dissociated in HSO^F. The following solvolysis equation is suggested: + C102S03F • C102 + S03F" The slight decrease in conductivities from the standard KSO^F for ClO^SO^F is due to a difference in the cation mobilities rather than an incomplete dissociation. The conductivity results for the solutes NOSO^F and NO^O^F lend support to this statement, as the + + observed order of conductivities at a given concentration, K > NO > + + N02 > C102 is in agreement with the expected trend for the cation siz The conductivity increase for (C102)2S30^Q is twice as large as found for C102S03F. The following mechanism of solvolysis is suggested + = (C102)2S3010 • 2C102 + S3010 S3°10" + HS03F " HS3°10 + S03F HS3°10" + HS°3F * H2S3°10 + S03F Overall Solvolysis: CC102>2S3°10 + 2HS°3F ~* 2C1°2+ + 2S03F" + H2S3°10 47 It is assumed that the acid H^S^O^ behaves as a non-electrolyte in HSO^F. This assumption receives some justification from studies on 95 the disulphuric acid solvent system , where H^S^O^ is reported to be essentially undissociated in H^S^O^, an acid of comparable strength to HSO^F. Further confirmation was obtained by solvolysis studies of ^S^O in HSO^F, showing similar conductivity values to those of (CIC^)2S30^Q at given concentrations. The conductivity curves for (CIO^^S^O^Q and the solutes discussed above are plotted in Figure 6. Conductivity measurements of (CIO^) ^^^^Q were also obtained in 100% H^SO^ to examine its solvolysis behavior in a weaker protonic solvent. The conductivity results along with the reference solute, KHSO^, are shown in Table 5 and plotted in Figure 7. (Here, y values are defined as the number of moles of hydrogen sulphate anions produced per mole of solute.) Dichloryl trisulphate dissolves slowly in 100% ^SO^ without apparent chemical reaction. Yellow to red solutions are obtained in the concentration range 0.005 molal to 0.040 molal, where (CIC^-^^^IO behaves as a strong base. Further additions of KHSO^ to the solution result in a continued increase in conductivity. Dichloryl trisulphate is found to be slightly more conducting than KHSO^. For the production of 1 mole of HSO^ in H^SO^ per mole of(C102)2S20^g, the following processes likely occur: 48 49 TABLE 5 (a): Specific Conductivities of (C102)2S 0 in H2S04 at 25.00°C. 2 2 2 Molality x 10 K x 10 102 Molality x 10 K X ro-l "I, (Sl cm ) (Sl cm ) 0.0000 1.045 2.263 1.204 0.0935 1.065 2.431 1.223 0.3192 1.073 2.511 1.234 0.5481 1.085 3.060 1.305 1.025 1.109 3.534 1.344 1.573 1.147 3.750 1.360 2.147 1.192 TABLE 5 (b) : Interpolated Specific Conductivities of (C102)2S3 and KHS04 in H"2S04 at 25.00°C. KHS04 (cio2)2s3 °10 2 2 Molality x 10 K_X 10^ K x 102 Y (Sl cm ) (Sl"lcm ) 0.000 1.043 1.043 0.500 1.053 1.081 1.03 1.000 1.068 1.109 1.04 1.500 1.097 1.143 1.04 2.000 1.139 1.183 1.04 2.500 1.180 1.227 1.03 3.000 1.240 1.279 1.03 3.500 1.289 1.339 1.04 4.000 1.356 1.408 1.04 50 FIGURE 7: I.OOl 1 1 —i -l r I 2 3 4 5. 2 1Q .Molality (cio2)2s3o10 2C102: + S3°10 S3°10_ + H2S04 " HS3°10 + HS04 The following equilibrium would account for the observed y values greater than 1.00: -y HS3°10 + H2S04 • t- H2S3°10 + HS04 As H^S^O.^ is a stronger acid than H^SO^, the equilibfiium exists far to the left. 3.3.3.3 1H NMR Spectra The proton magnetic resonance spectrum of a solution of ClO^O^F in HSO^F shows a single peak. The position of the singlet peak with respect to the solvent peak as external standard is shifted downfield with increasing solute concentration. This is in agreement 96 with the expected trend for the formation of fluorosulphate anions. Till results are. tabulated in Table 6 and plotted in Figure 8, along with the reference KSO^F values. The deviation of the ClO^O^F curve from linearity at mole fractions greater than 0.1 results from the imcomplete dissociation of ClO^O^F at high concentrations. FIGURE; s 52 H Chemical Shifts For Solutions of (CI02)OS02F in HSO3F o.o ( C!Op)OSOPF KS03F (Gillespie) 0.5 I.O 1.5 2.0 2.5 Mole Fraction x 10 1 1 TABLE 6 - H NMR Chemical Shifts of Solutions of C102S03F in HS03F * Mole Fraction C102S03F 6 (ppm) 0.191 1.367 0.108 0.900 0.035 0.350 0.017 0.167 * relative to HS03F external standard 3.3.3.4 UV-Visible Spectra The observed UV-visible spectra of solutions of C102S03F in HS0.F show three absorption peaks at X = 2250 A (e > 2500), o max max = 2700 A (e ~ 1950), and a shoulder of lower intensity at about Amax max ' 3 3300 A. The results are shown in Table 7, and the UV-visible spectrum, for the concentration 0.0159 M and a path length of 0.05 cm., is reproduced in Figure 9. T-he electronic spectrum of C102S03F in HS03F in HS03F can be + interpreted as essentially that of the chloronium cation, C102 . Assignments are made on the basis of the absorption spectrum of the 97 isoelectronic species S02 by Walsh. Walsh's diagram, illustrating the dependence of the binding energy of the molecular orbitals of a triatomic molecule on the apex angle, is reproduced in Figure 10. FIGURE 9 UV-visible spectrum of CI02S03F in HSO3F 4400 3800 3 200 Wavelength in TABLE 7 - Electronic Spectra of C102S03F in HSC^F (a) Observed Absorptions and Assignments 55 CIO, SO X (A) X (A) Assignment max t max 2250 >2500 2000 "400 2700 ~1'950 2940 B2 " 3300 sh ~ 900 3740 L Bl - (b) Extinction Coefficients for the Absorption at X = 2700 A max 10 Molarity Path Length (cm) Extinction Coefficient (e) 1.59 0.05 1960 1.80 0.05 1890 4.31 0.005 1950 6.33 0.005 1990 Average = 1948 FIGURE 10 56 90 I20 150 180° Angle BAB 57 The triatomic species ClO^ and ClO^ are reported to have apex angles of 110.5° and 115.3° respectively^. The removal of an electron II from the doubly occupied b^ molecular orbital of CIO2 results therefore in an increase of the O-Cl-O apex angle by about 5°. 11 (b^ is anti-bonding for Cl-0, but weakly bonding for 0-0). It is 11 then expected that the removal of the single electron from the b^ + M.O. in ClO^ might raise the apex angle to about 120° in C102 , o 99 which is close to that of the isoelectronic molecule S02> 119.5 In the ground state of C102+, the only low-lying vacant orbital it is b^ , and low energy electronic transitions will involve the excitation of an electron to this orbital. The two absorption bands at A = 3300 A and A = 2700 A are assigned to the allowed max max transitions, < *A^ and ^l" ^ne tnird transition ^2 -< *A^ in the group of low-energy transitions is forbidden. The allowed transitions are shown in Figure 10 and are labeled (1) and (2). The highly intense absorption band at X -2250 A involves 6 J r max higher energy transitions to and also to b^ from occupi ed 1 lower lying orbitals than the IT orbitals and a, A. • g Is. Thus the electronic spectra of solutions of C102S03F in HSO^F can be interpreted as arising from a triatomic species isoelectronic to S02- It is interesting to note that SO2 is colorless while C102+ is a red coloured species. As can be seen in Figure 9, the high intense bands of C102+ originate in the visible region at about 4400 A. The corresponding absorption bands for SO2 have been observed 97 at A = 3740 A and 2940 A , but they are of far lower intensity max J J fe for the stronger band is -400) and do not reach into the visible region, max 58 5.5.3.5 Migration Experiment Electrolysis of ClO^ solutions in a three compartment + diaphragm cell resulted in the migration of the red-coloured C102 species to the cathode, and the liberation of CIO,,. The following reaction occurs at the cathode: + C102 (red) + e y C102 (yellow) 3.3.3.6 Summary The following experiments give evidence for the existence of CIO,, in solution: 1. Conductivity measurements 2. NMR concentration-dependent chemical shifts 3. UV-visible spectra 4. Migration experiment + The first two experiments provide indirect evidence for C102 cations, as they show the formation of the corresponding fluorosulphate anions when C10,,S03F is dissolved in HSO^F. The possibility of a disproportionation reaction and the formation of chlorine dioxide were excluded by two further experiments: (a) ClO^SO^F could be recovered from .solution by evaporation of HSO^F and (b) A solution of C102S03F in HSC^F did not give an ESR signal. The electronic spectra of solutions of ClO^SO^F in HSO^F are consistent with the formation of ClO^* cations, while the migration experiment shows that the reduction of the species in solution results in chlorine dioxide. Satisfactory Raman spectra could not be obtained for solutions of the red-coloured species in HSO^F. Unfortunately, attempts to investigate the vibrational spectra of the red .chloryl compounds in the solid state were unsuccessful. High resolution Raman spectra could not be obtained due to the intense red colour of these compounds, while the high reactivity of the red chloryl compounds prevented the procurement of their infrared spectra! 60 CHAPTER FOUR CHLORYL HEXAFLUOROMETALLATES: CHLORYL COMPOUNDS (WHITE) PART II 4.1 INTRODUCTION A limited number of chloryl compounds have been known for some 11 18 time . They fall into two distinct classes. Chloryl fluoride, CIO2F , and its reaction products with strong Lewis acids such as BF^, AsF^. and SbF,.3'''"^5'''^ are colourless moderately reactive compounds. On the other hand, chloryl fluorosulphate, C1Q2S03F and dichloryl trisulphate, (C^^^S^O.^ are deep red-coloured solids at room temper• ature and are highly reactive compounds. Interestingly, all compounds + 3,100,102 have been postulated to contain the chloronium cation, CIO2 No structural investigations have been previously made, and no explanation has been put forth to account for the differences in reactivity and colour of the two classes of CIO2- containing compounds. It was shown in the last chapter that the chloronium cation, C102+, exists as a red-coloured species in a solution of fluoro• sulphuric acid when C102S03F and (CIO2)2S30.^ are used as solutes. The intense red colour and high reactivity in this latter class of compounds is then attributed to the presence of discrete C102+ cations. 61 It is then the purpose pf this chapter to resolve the apparent contradiction regarding the chloronium cation in the solid state by extending the solution studies performed in the last chapter to the colourless compounds, and ClC^AsFand by studying the vibrational spectra of CICLAsF, and ClO.SbF-. 2 6 2 6 4,2 EXPERIMENTAL Chloryl fluoride was prepared by the reaction of KCIO^ with 3 BrF^ in monel metal reactors (Figure 2) as according to Woolf . The purity of CIC^F was checked on the basis of the IR spectrum of the product. Arsenic pentafluoride and antimony pentafluoride were purified by trap to trap distillation. 3 Chloryl hexafluoroarsenate and -antimonate were prepared by reacting a very slight excess of chloryl fluoride with a known amount of pentafluoride in a monel metal reactor. The reaction was followed by weight changes. The ClO^AsF^ was purified by high vacuum sublimation. Solutions of ClO^F in fluorosulphuric acid were prepared by distilling purified chloryl fluoride in a metal vacuum line from a storage vessel into a weighing burette kept at -80°C. and containing a known quantity of doubly distilled HSO^F. This stock solution was used for addition to the acid in the conductivity cell. 62 4.3 RESULTS AND DISCUSSION 4.3.1 The Vibrational Spectra of C10„AsF. and C10„SbE^ z 6 2 6 The vibrational Raman and infrared absorption frequencies and intensities for ClC^AsF^. and ClO^SbF^ along with the frequencies and intensities for KAsF, and KSbF.. are listed in Table 8. The infrared 6 6 and Raman spectra of ClC^AsF^. are reproduced in Figures 11 and 12. The vibrational spectra of CK^AsF^ have also been reported by Christe 103 et al almost simultaneously to the publication of the research presented in this chapter. As will be shown, our interpretation of i the spectral results differs considerably from these authors, conclusions. Possible molecular structures for C10oMF,. where M = As or Sb are: z 6 + (a) an ionic solid C10_ MF ~ (b) a molecular 1:1 adduct C10oF-MFr with z 6 z 5 bridging over fluorine and (c) the same with bridging over oxygen. For model (a), the chloronium cation would be isoelectronic to sulphur dioxide, and the symmetry would be C^^. with three vibrational modes, both Raman and IR active. This is confirmed by the findings shown in Table 9. The vibration is similar to that of S02(s)"''^ which is split into two infrared and Raman active components by the crystalline field. It is interesting to note, however, that the average stretching frequencies for the CIO2 group (Table 10) are lower than that for SO2: = 1233 cm * v.s. ^>Q^Q + = 1166 cm * (IR) or 1168 cm * (Raman) — 2 -1 2 -1 in C10oAsF, and v„in + = 1174 cm (IR) and 1178 cm (Raman) for C10_SbF,. z o LIU2 z 6 TABLE 8: The Vibrational-Spectra of ClO^AsF, and CICLSbF OBSERVED WAVENUMBERS CcnT1) ClOJVsF^ ClO^SbF, KAsF, KSbF I.R. RAMAN I.R. RAMAN I.R. RAMAN I;R. RAMAJ1 1293 s 1296 w 1304 s 1313 w 698 s 692 s 655 s 661 s 1280 s 1281 w 1291 s 1295 w 382 br 580 ms 270 s 575 s 1080 w 1045 s 1050 ms 1053 s 375 ms 294 m 1045 s 725 w 810 vw 278 m ; 817 w 684 s 720 s ,br 706 s 725 s ,br • 572 s 680 s ,br 680 s 685 s ,br 518 m 670 s 563 s 370 s 640 s 649 s 515 s 585 ms 592 m •382 s 588 ms 372 s,sh 510 ms 518 w 490 ms 305 m 308 s 288 w 288 m,sh 270 s ,br 269 w TABLE 9: Assignment of the C102 Vibrations in CIO AsF, in C10oSbF, S02(s) Assignment C10, cio2 z o cio2 z o I.R. RAMAN I.R. RAMAN I.R. 1045 1045 1050 1053 1147 515 518 510 518 521 1293,1280 1296,1281 1304,1291 1314,1293 1330,1308 v3 (Bl} TABLE 10: C102 - Stretching Frequencies (cm -1) COMPOUND METHOD C102ASYM C102SYM C102 AVERAGE REFERENCE + Na C102 (s) Raman 832 804 818 This work , (aq. soln.) Raman 840 790 815 105 cio2 (g) IR 1111 943 1027 106 FC102 (1) Raman 1253 1097 1175 107 C102AsF6 (s) IR 1293,1280 1045 1166 This work ClOJVsF,, (s) Raman 1296,1281 1045 1168 This work C102SbF6 (s) IR 1304,1291 1050 1174 This work C102SbF6 (s) Raman 1314,1293 1053 1178 This work + K C103 (s) IR 960 910 935 108 (so2 IR 1330,1308 1147 1233 104} 65 67 On considering the increase in nuclear charge on chlorine, one would expect the VQ_Q + values to be higher. However, the infrared absorption frequency of the NO cation in a number of compounds has been found to vary from 2165 cm ^ to 2391 cm * ^. This variation has been attributed to anion-cation interaction which results in the lowering of the frequency of the N-0 vibration. For a corresponding MF^ anion with 0^ symmetry, two infrared F are active vibrations, (F' ) and ( iu) expected together with three Raman active vibrations, v, (A, ), v~ (E ) and vr (F„ 1. 1 lg' 2 g' 5 2gJ The observed infrared spectrum of C102AsF^ is more complex than expected. In the region of v^, a strong absorption at 685 cm ^ occurs with a shoulder at 725 cm ^; this is in contrast to the spectrum of KAsF^ where a single peak is observed for at 698 cm ^. (A similar observation has been made by Clark and O'Brien for the infrared spectrum 109 of (CH^^SnAsF^ , where the AsF^ group is said to act as a bridging group and the symmetry appears to be lowered to In the region of v^, a doublet is observed at 382 and 372 cm ^. In addition, a strong absorption peak at 563 cm ^ is observed, but is not found in the IR spectrum of C102SbF^. It must therefore be due to the.AsF^ group. The Raman spectrum of C102AsF^ also shows more absorption peaks than expected. The positions of ," and v,. at 684, 572 and 370 cm ^ respectively are at lower frequencies than they are observed in KAsF,. 68 From the position of the expected mode, v^, at. 572 cm it is concluded that it is identical with the observed 563 cm * vibration in the infrared spectrum. The 725 cm * vibration is also weakly Raman active, and the possibility that this vibration is a combination mode seems rather unlikely. From the findings of the vibrational spectra, it seems that the symmetry of the AsF^ group isilowered considerably from the expected 0^ symmetry, resulting in the relaxation of the mutual exclusion rule for IR and Raman active vibrations and a general shift of all frequencies in comparison with other AsF^. compounds. The tentative assignment of the observed frequencies is shown in Table 11. The reason for the observed symmetry lowering of the anion can be seen in appreciable anion-cation interaction via fluorine bridges, already reflected in the shifts of the Ci02 stretching vibrations (Table 10). It is important to note that a possible anion-cation polarisation should not result in the observed trend for v^,^ and would be expected to be larger for small and polarising cations such as N0+. This expectation is not confirmed by experiment^? The vibrational spectra of the corresponding hexafluoroantimonate compound are even more complex. No certain assignments for the SbF, anion can be made; crude assignments are given in Table 12. The infrared spectrum of ClO^SbF^ reveals a broad absorption peak in the region of with components at 640, 670,680,720 cm *. The peaks at 640, 670 and 720 cm * are assigned tentatively to the components of v^, while the absorption at 680 cm * is believed to be due to the infrared forbidden v^. 69 TABLE 11: Assignment of the AsF, Vibrations 6 AsF6" in C102AsF6 AsF6 in KAsF6 Assignment AsF6~ (0h) I.R. RAMAN I.R. RAMAN (685)* 684 - 692 v, (A ) 1 lg' (563) 572 - 580 v2 (E ) 725 (725) 698 - v7 (F, ) 3 lu 382 - 382 - v, (F. ) 4 v lu' (372) 370 - 375 v (F ) J zg * Frequencies in parenthese are forbidden for AsF^ with 0^ symmetry TABLE 12: Assignment of the SbF ~ Vibrations 6 SbF6" in C102SbF6 SbF6" in KSbF^ Assignment SbF6" (0h) I.R. RAMAN I.R. RAMAN 680 680 - 661 v, (A ) 1 lg' ] - ^ v2 ^ 586 592 -j - 575 v0 (EJ 51 720 n 706 n 655 - V (F ) 11 3 lu 670 I 649 ] 640 270 269 270 - v, (F, ) 4 v lu' 305 n 308 -i - 294 -j v5 (F2 ) 288 -» 288 J 278 J V6 70 A broad absorption is also observed in the region of v^, which is assigned at 270 cm *; peaks at about 305 cm ^ and 288 cm ^ are assigned to vrJ but lack of resolution and precision in this range does not differentiate these peaks from components of v^. In addition, is observed in-:the infrared spectrum as a medium strong band at 586 cm The Raman spectrum of C102SbFg shows a similar complexity. Strong Raman active bands are found at 706, 680 and 649 cm * in the region of v,, while v„ is observed as a doublet at 592 and 588 cm ^. & v 1 2 Vj. is observed at 308 cm * with a shoulder at 288 cm * . The loss of octahedral symmetry for the SbF^ anion in C102SbF^ is quite apparent. The splitting of the degenerate vibrational modes, the relaxation of the mutual exclusion rule, and the absorption peaks assigned to the SbF^ group appearing at relatively high frequencies are all features of the vibrational spectra which can arise from anion-cation interaction. Such an interaction, which is more pronounced for C102SbF^ than C102AsF^, probably occurs via fluorine bridges. The average CIO2- stretching frequencies for a number of ei02~ containing compounds, along with the values for SO2 are shown in Table 10. It is noteworthy that for the hexafluoroanion complexes, the average CIO2 stretching frequency in each case is about the same as that for chloryl fluoride where v + = 1175 cm ^. 71 Th e bonding in CIC^F involves a short and strong chlorine- oxygen bond and a weak and long chlorine-fluorine bond as is evident 107 from the reported force constants . This is reminiscent of the situation in FNO1*''' and FI^*"''2 as well as for ®_^2^^ anc^ ®2^^ where the observed element fluorine bonds are found to be longer than the normal single bonds. Consequently, chloryl fluoride is a highly polar molecule. The Cl-F bond has been interpreted as the interaction between a singly occupied 2p orbital of fluorine to an anti-bonding TT* M.O. of chlorine dioxide, resulting in a highly delocalized multi-centre bonding. 114 (This type of bonding,labelled by Spratley and Pimentel as 113 {ir*-p} a bonding,was used by Lipscomb in the description of 0_^2' where the 0-F bonds result from the interaction of a high lying singly occupied antibonding TT* orbital of 0^ and a singly occupied 2p orbital on fluorine.) It seems possible that this {7T*-p} a interaction is retained when CIO2F reacts with AsF^ and SbF,. , causing no alteration in the Cl-0 bond order as evidenced by the vibrational spectra. It was shown in the last chapter that the electronic transitions responsible for the red colour of the C102+ cation involve the low- ii lying b^ anti-bonding molecular orbital. If, however, this orbital became partly occupied as a result of a {Tr*-p} a interaction involving 11 the b^ M.O. and a 2p orbital of one of the flourine atoms of the anion, the electronic transitions would not be able to occur, and the compound would very likely appear colourless. 72 It is feasible that this is the case with C102F and its reaction products with Lewis acids. In the case of C10_AsF^ and C10„SbF., I o I o this interaction is not necessarily restricted to one fluorine, and two fluorine atoms may be involved. The {ir*-p} a interaction in white chloryl compounds provides a good model to explain their lack of colour and to rationalize the average CIO,, stretching frequencies of C102AsF^ and CIO SbF^',* but other electronic interactions probably exist between the anion and the cation. In summary, the vibrational spectra show the white chloryl compounds ClO^AsFg and ClO^SbF^ to be predominantly ionic, but with strong interaction between the anion and the cation via fluorine bridging. This anion-cation interaction lowers the charge and reactivity of the C102 group and accounts for the polyfluoroanion complexes and C102F occurring as colourless compounds. 103 In contrast to these conclusions, Christe et al have interpreted the structure of ClO^sF^., on the basis of the same vibrational spectra, as consisting simply of discrete ions and have calculated force constants + for C102 . In our opinion there seems to be strong evidence for fluorine-bridging in ClO^sF^ and related compounds, and .elaborate force constant calculations can not be justified. + As shown in chapter 3, the C102 cation exists as a red-coloured species in HSO^F. It now became interesting to examine the behavior of the white chloryl compounds in HSO^F. 4.3.2 Solution Studies The solutes KAsF, and CICLAsF, are soluble and behave as bases 6 2 6 in HSO^F as indicated by the conductimetric measurements. Specific conductivities at various concentrations are listed in Table 13 along with the reference KSO^F values. Chloryl hexafluoroarsenate dissolves in HSO^F to form a red-coloured solution, the colour being characteristic + for the C102 cation, while the solute KAsF^ gives rise to a colourless solution. However, whereas KAsF^ is only moderately soluble, the 2 solubility limit being 5.0 x 10~ molal at 25.0° C, the C102~ compound is soluble to a far greater degree, up to 1 - 2 molal. The specific conductivity v.s. molality plots are shown in Figure 13. The specific conductivities of the ClO^AsF^ solutions are slightly higher than the ones for KAsF^, whereas the situation was reversed when C102S03F and KSO^F were compared as solutes. These differences in solubility and conductivity are unexpected. The role of the AsF^ anion in HSO^F was not investigated as the concentration ranges were 19 too low for meaningful Raman and F NMR studies. Possibly, as indicated by the y values, the AsF^ anion undergoes disproportionation or chemical exchange in HSO^F: As this problem was not resolved and to simplify the solution studies, greater attention was focussed on the solution behavior of C10„F in HS0,F. TABLE 13: (a) Specific Conductivities in HS03F at 25.00°C. ClOJVsF, KAsF. C10oF Z o o Z Molality x It)2 K x 104 Molality x 102 K X 104 Molality x 102 K X 104 (ft 1 cm (ft '''cm "*") (ft "'"cm ^) 0. 0000 1. ,236 0..00 0 1. .127 0..000 0 1. ,128 0..21 8 3. ,182 0,.20 2 5, .504 0..170 0 3. .665 0.,52 3 6.,71 2 0.,37 2 8..55 8 0..469 0 11. ,94 1. 248 17. ,89 0..55 4 10. .81 0..595 4 15. .40 1. 972 26. 44 0.,87 9 14. .41 0.,782 0 20. ,20 2. 785 34. ,17 1. ,306 18. .93 0.,858 7 22. ,02 3. .964 42. ,93 1. .804 23, .21 1. .1205 30. .82 5. ,025 50. ,04 2, .147 26, .05 1. .482 43. .47 6.,30 9 58. ,38 2. .863 31, .74 2. ,094 63. ,04 3. .403 36, .29 2. ,611 78. ,87 4. ,119 41. .78 3. ,035 91. ,71 4. .563 45, .78 3. .759 112. ,5 4. .892 143. .3 6,.13 4 175. ,8 TABLE 13: (b) Interpolated Specific Conductivities KS0,F KAsF ClO^AsF, C10oF 3 6 2 6 2 MOLALITY K x 104 K x 104 K x 104 K X 104 x IO2 ( ft"1™"1) (fi_1cm_1) (^"1cm"1) (fi"1cm~1) 0.00 1.085 1.127 1.236 1.329 0.25 7.0 6.8 4.0 7.0 0.50 13.6 9.8 6.5 13.2 0.75 19.7 12.3 13.2 19.5 1.00 25.8 14.5 16.2 26.1 1.50 38.0 21.7 22.5 44.1 2.00 50.0 24.3 27.3 61.0 2.50 62.5 27.7 31.8 75.8 3.00 72.7 33.3 36.0 90.2 3.50 84.9 37.3 40.0 104.7 4.00 97.0 40.5 43.6 118.9 4.50 106.8 45.1 47.3 133.0 I I 1 1 1 1 n , 0.5 I.O 2.0 3.0 4.0 5.0 6.0 T-O IO2 Molality Fig. 13 Solutions of Clt^F in HSO^F are red-coloured and highly conducting (Table and Figure 13). The UV-visible absorption spectrum corresponds + 19 to that of C102 in HSO^F. A F NMR spectrum of the solution reveals the presence of HF as the only fluorine-containing species besides the solvent. The Y values are higher than the reference KSO^F values, which is not unexpected since KF has also been found to be 45 more conducting than KSO^F (because HF behaves as a weak base in HSO^F). In addition, the preparation of the stock solution of C102F in a pyrex container may have resulted in side reactions and a very slightly higher conductivity value. At concentrations below _2 10 molal,Y values slightly less than 1 are obtained. This is very likely due to the presence of SO^ as an impurity in the HSO^F 45 which behaves as an acid in this solvent system . Chloryl fluoride appears to ionize in HSO^F as follows: + C102F + HS03F • C102 + HF + S03F~ HF + HSO.F »• H_F+ + S0,F~ This would result in y values greater than 1.0. The-described behavior of C102F in HS03F was studied previously by Schmeisser and Fink*^''' who postulated the equilibrium: C102F + HS03F < HF + C102S03F It is reasonable to assume that in dilute solutions, HF and C102S03F will be ionized and the equilibrium shifted to the right. 78 an In summary, C102F d ClC^AsFg produce red-coloured solutions in HSO F as the result of the formation of discrete CIO cations. In addition, C10„SbF, dissolves in HS0„F to form a red-coloured solution. ZD o It seems then that the fluorine bridge bonds are broken when the white chloryl compounds are dissolved in the strongly protonic solvent of high dielectric constant: Cl0 + C102F (colourless) • 2 (red) + F~ C10oAsF, (white) >- CIO* (red) + AsF 2 6 <•""—-' K J 6 nSbF^ (white) U- CIO* (red) + SbF," 2 o z o The above reactions indicate the first step in the ionization + of these compounds in HSO^F. Discrete C102 cations are formed and stabilized by solvation with HSO^F. CHAPTER FIVE DICHLORYL HEXAFLUOROSTANNATE : CHLORYL COMPOUNDS (WHITE) PART II CONTINUED 5.1 INTRODUCTION It became interesting to extend the study of chloryl compounds to dichloryl hexafluorostannate, (C102)2SnF^. In addition to vibrational spectroscopy, it would be now possible 119 to use Sn Mossbauer spectroscopy as a means of obtaining structural information and detecting possible anion-cation interaction in this compound. The distortion of the anion octahedron as a result of anion- cation interaction should result in a non-vanishing electric field gradient and subsequently in a quadrupole splitting. An asymmetric distribution of electron density in valence orbitals about the tin nucleus will cause a quadrupole splitting. Although Greenwood and Ruddick''''''^ first interpreted the quadrupole splitting in hexacoordinated 117-118 tin as a pit-dir bonding effect, it is now generally believed that a-bonding effects, that is , differences in bond polarities, are the dominant factor in producing a quadrupole splitting. 119 It has been shown on adducts of the type SnCl^" 2R2SO that molecular distortion caused by bulky ligand groups can give rise to small quadrupole splittings. 80 After the research described in this chapter had been completed, 120 121 we learned of a publication by Sukhovrekhov et al ' ,who have found quadrupole splitting for SnF^ complexes with heterocations BrF2+, BrF^*, CIF^* and IF^+. Unfortunately, the vibrational spectra of these complexes have not yet been reported. 119 The Sn chemical shift is commonly used as a measure of the 5s electron density around tin. It has been previously reported that 122 the hexafluorostannate anion, SnF^ , has a very low chemical shift The reported value both for I^SnF^ and Cs^SnF^ of -0.44 mm/sec is 4+ 123 even lower than the calculated value for a Sn ion. Therefore, electronic interactions between the SnF^ anion and heterocations o should affect the chemical shift value. Previous reports on the vibrational spectrum of the hexafluoro- 124—126 stannate anion are both incomplete and inconsistent . Consequently, this chapter is also concerned with the complete assignments of the fundamental vibrations of the SnF, anion. b 5.2 EXPERIMENTAL 3 Chloryl fluoride,CIO2F, was prepared as according to Woolf . For the preparation of (CIO2)2SnF^, the method of Schmeisser and Fink""^ was adopted but with modifications and without the use of CFClg as a solvent. The preparation is described as follows: To minimize the hazard involved in handling CIO2J 89.1 mmoles of CIO2F were introduced in small portions into a monel metal reaction vessel (Figure 2) containing 8.43 mmoles of purified SnCl^ (Fisher Scientific Co.) by vacuum distillation in a metal line. The reactor was warmed to room temperature and the volatile by-products CIC^ and Cl^ were removed by pumping. This procedure was repeated several times until all the ClO^F was added. The volatile products were removed again by pumping and the involatile solid, indicated by the weight increase, corresponded to (ClO^^nF^. Potassium hexafluorostannate was purified by recrystallization from anhydrous HF in a polyethylene reactor. Sodium Hexafluorostannate was prepared from Na2Sn03 (Fisher Scientific Co.) in 49% HF (Fisher Scientific Co.). The infrared spectrum of K^SnF^ was recorded on a Perkin Elmer 301 High Resolution Grating Spectrophotometer for the range 666-166 cm * Nujol as a mulling agent and Csl cell windows were used. 5.3 RESULTS AND DISCUSSION 5.3.1 Vibrational Spectra The infrared and Raman data for the compounds Na„SnF,, K„SnF, z 6 z 6 and (C102)2SnF£ are listed in Table 14. The vibrational frequencies for the SnF ~ anion in K„SnF^ and Na„SnF, are tabulated in Table 15. 6 2 6 2 6 The previously reported values for Na2SnF^ and K^SnF^, as well as a number of other SnF^ compounds,are included. The alkali metal compounds, I^SnF^ and Cs^nF^ are reported to have the trigonal I^GeF^ structure with the anion having D^^ site 127 128 symmetry. ' An orthorhombic structure with a D^ symmetry site 127 for the anion is found for Na0SnF,. z o TABLE 14: Vibrational Frequencies For SnF Compounds (cm ) Na^SnF, K2SnF Ccio2)2 SnF£ 6 2 0 0 I.R. RAMAN I.R. RAMAN I.R. RAMAN 590 sh 594 vs 582 w,sh 628 vw,sh 1302 vs 1309 m 560,vs 480 m 557 vs 598 vs 1290 vs 1293 m 475 vw,sh 416 w 480 vw 478 vs 1076 s 1080 vs 367 vw 420 vw 423 w 1072 s 255 ms 367 vw 368 w 634 s,sh 628 m-: 162 m 257 s 258 s 624 s,sh 254 s, sh 606 s 613 s 163 m 561 vs 563 vw 541 s,sh 549 s 522 s,sh 519 m 470 m-s 482 m 465 •W,: 423 w 314 w 367 w vw = very weak; s = strong; w = weak; vs = very strong; m-= medium; sh = shoulder oo TABLE 15: Vibrational Frequencies for the SnF, Anion Compound v,(A ) v„(E ) v-CF, ) v. (F, ) vc(F„ ) Reference r 1 lg 2 g 3 lu 4 lu 5 2g K2SnF6(s) 598 478* 582,557 257,254 257 This Na2SnF6(s) 594 480* 590,560 255 This (NO)2SnF6(s) 593 481* 585,555 258 256 29 (N02)2SnF6(s) 590 478 590,565 29 Cs2SnF6 (s) 572 460 577,555 256,239 247 126 K2SnF6.H20(s) 593 620 564 342 125 Na2SnF6(s) 592 477 559 300 252 129 (NH4)2SnF6 (aq. soln.) 585 470 241 126 * denotes both IR and Raman active 130 As a consequence , the infrared active vibrations ^(F^) and v^(F^u) for K^SnF^ should both be split into the two components a^u and e^. A similar splitting is expected for Vj-t^ ) resulting in the components a^ and e . The high resolution infrared spectrum of I^SnF^ is reproduced in Figure 14. As can be seen, both at cm and at 257 cm * have shoulders at 582 cm * and 254 cm * respectively. In addition, v2(E ) appears as a weak absorption peak at 478 cm *. The Raman active vibration v,., however, is found as a single peak in the Raman -1 126 spectrum (at 257 cm ) in agreement with previous work for C^SnF^. 125 -1 Kriegsmann and Kessler have reported a shoulder at 620 cm in the infrared spectrum of K^SnF^.H^O, and have assigned the same to the Raman active vibration which becomes weakly active in the 126 infrared. As remarked by Dean and Evans and in view of the values reported for for all other SnF^~ compounds, this assignment is undoubtedly incorrect. The previously reported vibrational spectra of Na^SnF^ by Begun 129 and Rutenberg are confirmed except for their assignment of at 300 cm * and their failure to report a splitting for v^. No unusual features are observed for the Raman spectra of the SnF^ compounds listed in Table 15. The spectra are very similar 126 to that obtained for an aqueous solution of (NH^^SnF^ . In the infrared spectra, the splitting of the and modes and the occurrence of the forbidden transition are attributed to site symmetry effects. 86 However, no simple explanation based on an ionic compound with an anion site of lower symmetry than 0^ appears to be possible for (ClC^^SnF^. Its vibrational spectra are reproduced in Figures 15 and 16. The high reactivity of this compound relative to the corresponding alkali metal complexes prevented high resolution infrared spectroscopy. Also, attempts to observe low frequency Raman bands failed. Consequently, the interpretation of the vibrational spectra of (ClG^^SnF^ is restricted to the range down to 300 cm ^ and the anion fundamentals v±>v2 and V3' Focussing our attention to the anion as a sensor for anion-cation interaction, the following features of the Raman spectrum, shown in Figure 16 in the range of 800-300 cm * are noteworthy: v^, which is split into 2 components, presumably as a result of solid-state splitting, is shifted to higher frequencies as it occurs in K^SnF^, and is observed at 628 and 606 cm The normally Raman inactive is observed as a strong absorption band at 549 cm * with a weak shoulder at 563 cm ^. occurs as a weak absorption band at 470 cm *. Absorption bands at 423 and 367 cm * appearing as very weak and diffuse peaks are seen in all SnF^ Raman spectra and are due to the pyrex tubes. (For Na^SnF^ and K^SnF^, the moderately strong band at about 165 cm also results by background absorption from the pyrex tube. FIGURE 15: The Infrared Spectrum of (C10o)oSnF/. 2 2 6 (CI02)2SnrJ FIGURE 16: The Raman Spectrum of (CIO ) SnF The infrared spectrum of (ClO^^nF^ shown in Figure 15 has a very broad, barely resolved band in the region from 630 cm 1 to 520 cm 1 Assignments for the anion vibrations gre difficult for two reasons: (1) Resolution is rather poor and (2) it becomes difficult to distinguish between a true splitting or additional peaks caused by a breakdown of the mutual exclusion rules. Two components of are assigned to the absorption at 561 cm 1 and a shoulder peak at -1 -1 541 cm . v occurs as a forbidden transition split at 634 cm and 606 cm ^, while is observed as a medium strong band at 478 cm 1. It is difficult to reconcile the vibrational spectra of (C10_)oSnF, I 2 o with an ionic structure and anion site of lower symmetry than octahedral As in the case of the spectra of dO^sF^ and ClO^SbF^, strong anion- cation interaction is suggested. Anion-cation interaction should affect the vibrational modes of the cation as well. The absorption peaks in the spectra due to ClO^* are shown in Table 16: TABLE 16: Vibrational Frequencies of CIO + in (CIO ) SnF (cm"1) I.R. .... RAMAN ASSIGNMENT 1302,1290 1309,1293 v3 C102 Asym. 1076,1072 1080 v C102 Sym. 522 519 v2 C102 Bending As ClO^ is isoelectronic with SO^j it is reasonable to expect C2v symmetry. This would make all three vibrations non-degenerate and both infrared and Raman active. A possible interaction should therefore affect only the band positions. This is illustrated by + a comparison of the C102 frequencies for the white chloryl compounds. The average stretching frequencies are raised by 5-10 cm * when going from C102AsF6 over C102SbF6 to (ClO^SnF^ In particular, the symmetric vibration for the hexafluorostannate anion is shifted upwards by about 30 wave numbers compared to C102AsFg and ClO^bF^. Again, interaction between the anion and the cation will likely occur over fluorine bridges. In the last chapter, the involvement of the b1 molecular orbital (of C102 ) in bonding in C102F, C102AsF6 + and C10„SbF£ was invoked to explain the absence of colour for CIO . 26 2 + The noted increase in the average stretching frequencies of the C102 cation in (ClO^^nF^ can be explained by electron transfer possibly involving the b^" and the a^ g molecular orbitals on C102 , where electron withdrawl from the cation results in an increased electron density in the tin-fluorine bonding region. If this idea is correct and provided an electron transfer stabilizes certain Sn-F bonds over others, the effect on the Mossbauer spectrum should be noticeable. 91 5.3.2 Mossbauer Spectra The Mossbauer data for the SnF6~ complexes are listed in Table 17 for 80° and 298° K. The accuracy limit for the Mossbauer data is judged to be - 0.03 mm/sec for both the isomer shift and the quadrupole splitting. The linewidth values are also listed in Table 17, as well 131 as the room temperature effect R , where R is defined as eonoo/eono, 298 80 e being the magnitude of the Mossbauer effect. All compounds studied, 29 as well as (N0)_SnF, and (N0_) SnF^ included in the table for ZD 2 2 ° comparison, gave well-resolved spectra at room temperature, a feature 131 132 49 normally associated with intermolecular association ' ' and polymeric structures. As indicated by the vibrational spectra, the exceptional position n of (C102)2S FD is also found for the Mossbauer spectrum. The isomer shift is the highest reported in this series, indicating an increase in s-electron density around tin. As shown in Figure 17, a well-resolved quadrupole splitting of about 1.00 mm/sec is observed. The A value is comparable to the previously reported values for (Eh^^SnF^ i on and CBrF4)2SnF6 of 0.80 and 1.15 mm/sec respectively . The occurrence of quadrupole splitting in these compounds indicates considerable departure from octahedral symmetry for the SnF^- anion, which is best explained by anion-cation interaction. In contrast, the Mossbauer spectra of K„SnF., Na.SnF,, (N0)_SnF. 2 O 2 O ZD and (ITC^^SnF^ all show broad singlets. This line broadening is attributed to unresolved quadrupole splittings which can be produced as the result of the low site symmetry of the anion. 92 It is possible by means of computer fitting to resolve this splitting. This is the case for the Mossbauer spectrum of (NCO.SnF,.. 2 2 b In conclusion, the vibrational and Mossbauer spectra of (C10_)„SnF, 2 2 b are inconsistent with a simple ionic structure for this compound. Allowance must be made for a substantial interaction between the anion and the cation. (Similar conclusions for the fluoro-halogen heterocation complexes (CIFJ „SnF,, (BrFJ-SnF., (BrF,)„SnF^ and 2 2b 22b 42b 120 121 (IF^^SnF^ have been previously made by Sukhovrekhov and Dzevitskii ' on the basis of Mossbauer spectroscopy.) It appears, however, not possible to discriminate between and site symmetry for the tin atom, as the two most obvious possibilities for a distorted octahedral anion, mainly because no satisfactory vibrational spectra in the region of v. and vr could be obtained. 5 4 5 TABLE 17: Mossbauer Data For SnF Compounds Compound Temp. Isomer Shift Quadrupole Splitting Linewidth R (°K) 6 (mm/sec) A (mm/sec) Na^SnF. 80 -.480 1.77 2 6 0.71 298 -.528 - 1.35 K SnF, 80 -.432 1.59 2 6 0.61 298 -.468 - 1.11 (NO)2SnF6 80 -.421 - 1.45 298 -.456 - 1.10 (N02)2SnF6 80 -.431 .769 1.46 ' 1.57 0.58 298 -.504 .680 1.11 1.75 (C102)2SnF6 80 -.403 1.008 1.46 1.46 0.71 298 -.433 .956 1.16 1.11 94 CHANNELS FIGURE 17: The Mossbauer Spectrum of (C10J-,SnF CHAPTER SIX STRUCTURAL STUDIES OF HEXAFLUOROARSENATES AND -ANTIMONATES OF FLUORO-HALOGEN HETEROCATIONS 6.1 INTRODUCTION Hexafluoroarsenate and -antimonate complexes of fluoro- and oxy- 133 element heterocations have been known for some time . These complexes are formed by the reaction of the Lewis acids AsF,. or SbF,. with an element fluoride or element oxyfluoride to form a 1:1 complex: EF + MFr EF ,MF. n 5 n-1 6 EO F + MFr >- EO ,MF^ m 5 m-l 6 E = P, S, Se, Cl, Br, I and others M = As, Sb A large number of fluorides can act as Lewis acids in the above equations (e.g. BF^, SiF4, GeF^, TiF^, VF5, NbF5, PF5 and others). The two most likely and extreme forms for the molecular structure of these complexes are: (1) a covalent structure, e.g. EF^MFj., a 1:1 adduct involving bridging over fluorine; (2) an ionic structure consisting of discrete ions, where the Lewis acid has completely extracted a fluoride ion from the element fluoride or oxyfluoride. 96 A number of such hexafluoroarsenate and -antimonate complexes are shown in Table 18. These complexes have been suggested to be 3 ionic in the past. Woolf , for example, postulated ionic strucures for ClOJVsF, and C10_SbF,. On the basis of conductivity data in z o z o 134 liquid frrf^j Woolf and Emeleus proposed for the 1:1 complex produced from the reaction of BrF„ and SbF_ the structure BrF_+SbF^ 3 5 z 6 in lieu of BrF^SbF. . Recently, the structural aspects of this class of compounds 133 have received much attention . However, there has been some controversy in the structural interpretation. The vibrational 103 135 136 spectra of the related compounds C102AsF^ and ClF2AsF^ ' have been interpreted using a strictly ionic model as evidenced by elaborate force constant calculations. On the other hand, evidence for anion-cation interaction in C102AsF^ was presented in Chapter 4. A similar interpretation is indicated by the x-ray 137 diffraction study of BrF2SbF^ , where a distorted octahedral -SbF^ unit is found, exerting fluorine bridge bonding to bromine in the cation. The purpose of this chapter, is to extend the study of the chloryl compounds C102AsF^ and C102SbF^ (chapter 4) to the related compounds of fluoro-halogen heterocations. In particular, a comparative study of ClF2AsF^, ClF2SbF^ and BrF2SbF^ is made using infrared spectroscopy down to 250 cm * and laser Raman spectroscopy. The cations C1F2+ and BrF2+ are chosen because they are triatomic, presumably bent and of comparable size to the C102+ cation. TABLE 18: AsF,~ AND SbF," COMPLEXES OF FLUORO- AND OXY-ELEMENT HETEROCATIONS PERIOD GROUP 5A GROUP 6A GROUP 7A N2FAsF6 N0AsF6 2 N2F3AsF6 N02ASF6 NF.AsF, NOSbF, 4 6 6 FoN0AsF, NO„SbF, I o I o PF2AsF6 SF3AsF6 C£F2AsF6 3 PF2SbF6 SF3SbF6 C£F2SbF6 C£02AsF6 CJl02SbF6 SeF3AsF6 4 SeF3SbF6 BrF~SbF, l o to From a study of a number of related nitrogen -oxy and -fluoro hexafluoroarsenates and -antimonates11^, a narrow frequency range for the active vibrational fundamentals of AsF, and SbF. has been 6 6 established. These compounds were suggested to have ionic structures, and therefore should be useful in comparison with the vibrational spectra of the corresponding -AsF^ and -SbF^ complexes of fluoro- halogen heterocations. Anion-cation interaction in the latter compounds is expected to affect the vibrational modes of the highly symmetrical anion in the following ways: (1) removal of the degeneracy of the E and F modes; (2) relaxation of the mutual exclusion rule for Raman and IR active vibrations and (3) a general shift in frequencies from where they occur in the vibrational spectra of the corresponding S an( ionic potassium salts, KA FQ i KSbF^. As according to Walsh's 97 predictions , the cations are all expected to have symmetry, where all three vibrations are non-degenerate and both Raman and infrared active. Shortly after the time that this study had been completed, 138 Gillespie and Morton reported the Raman spectra of ClF^AsF^ 13S 136 and ClF^SbF^. In contrast to earlier work ' , they have concluded that there is fluorine bridging between the CIF^* cation and the anions. 99 6.2 EXPERIMENTAL' All compounds were found to be very sensitive towards moisture and highly reactive. Either the monel metal vacuum line or dry-box was employed in the transfer of materials. The 2-part monel metal reactions vessels (Figure 2) were used. All reagents were purified by distillation. Following purification, SbF^ was added to the reaction vessel in the dry box, while AsF,. was distilled directly in the vacuum line. Except for the use of monel metal vacuum line equipment and reactors, the reported preparative 139 140 routes were employed ' . The reactions were followed by weight, with the most volatile component in excess. 6.3 RESULTS AND DISCUSSION 6.3.1 Results The infrared and Raman vibrational spectra of ClF^AsF^, ClF„SbF, and BrF„SbF, are reproduced in Figures 18 - 21, while z 6 z 6 the frequencies and intensities are listed in Table 19. The Raman spectrum of ClO^SbF^ is included in Figure 21 with the ~SbF^ complexes for the purpose of comparison. Again, as in Chapters 4 and 5, assignments can be made on the basis of predominantly ionic structures, and with reference to the ionic salts KAsF, and KSbF,. 100 TABLE 19: The Vibrational Spectra of ClF2SbF6, ClF_AsF. and BrF„SbF, ZD ZD ClF„SbF. ClF.AsF. BrF„SbF. 2 6 2 6 2 6 IR RAMAN IR RAMAN IR RAMAN (cm 1) (cm 1) (cm 1) (cm (cm 1) (cm-1) 1310 vw 1064 vw 830 sh,vw 826 sh, s 2315 ms 706 vs 1295 vw 1040 vw 813 br ,w .809 vs 673 vs 693 w 1065 vw 833 m 688 vs 693 vs 663 vs ,br 679 s 830 s 810 vs 605 m 603 w 638 s 641 s 803 s 662 vs 560 wm 544 s 535 m, sh 552 vs 677 s 644 w 520 w 382 s 523 ms 523 m 660 s 596 w 383 vs 372 s 495 ms 493 ms 647 s 542 mw 310 w 370 w,sh 367 m 604 s, sh 537 mw 285 m,sh 284 m,sh 582 mw 385 mw 270 m,sh 270 m 375 mw 292 w 265 m 290 w,sh 282 m 283 w,sh 267 m 265 s Vw = very weak, mw = medium weak, w = weak, m = medium, s = strong vs = very strong, sh = shoulder, br = broad —| ' 1 1 —I r- lOOO 800 600 400 250 Wave number (cmr1) FIGURE 18: The IR Spectrum of ClF2AsF6 FIGURE 19: The Raman Spectrum of ClF-AsF, 2 6 CIE; Asf^ 900 800 700 600 , 500 400 300 200 Wavenumber (cmH ) The Raman Spectra of C\FZ SDrf , BrffSbrj and CI02Sbf^ r i i i 1 1 1— 900 80Wavenumbe0 700 600 50r 0(cm 400-1) 300 FIGURE 21 105 138 Agreement with the results of Gillespie and Morton is good except for minor differences. A small splitting of v^ClF^* observed by these authors in the Raman spectrum of ClF^SbF^ and interpreted as a solid state splitting was not resolved in this work. In addition, a shoulder of weak intensity at 644 cm 1 was observed. 6.3.2 Assignment of the Cation Vibrations The vibrational frequencies assigned for the cationic groups CIF^* and BrF^* are presented in Table 20. The vibrational modes of + the C1F2 cation can be assigned on the basis of comparison with the isoelectronic species ClO^ , which has v^, \>2, and at 804, 432 1 + and 832 cm respectively. The assignments for C1F2 differ from 136 those previously reported by Christe and Sawodny in that the + 1 C1F2 bending mode, v^, is assigned to the band at 382 cm (in C1E> AsF,) rather than 544 cm 1. In addition, these authors were unable 2D to resolve the CIF^* asymmetric and symmetric stretching frequencies in the Raman spectrum which in this work were found noticeably separated at 826 and 809 cm 1 respectively. + -1 -1 The reassignment of V2C1F2 (382 cm v.s. 544 cm ) is made for the following reasons: (1) The difficulty in assigning the bending mode of C1F* in ClFJ\.sF, lies in the fact that the AsF ~ Z Z O O anion also absorbs in this region. However, the vibrational spectra 1 of ClF0SbF^ also show a band near 380 cm , where the SbF^ anion 2D 6 does not absorb. (2) The value for of 382 cm 1 is in much closer agreement to that for the isoelectronic species ClO^ where v2C102 = 432 cm ^, 106 TABLE 20 Vibrational Frequencies of the Species C1F + and BrF„+ (cm *) Species Compound Method v. ,,v„ v + C1F2 ClF0AsF£ Raman 809 382 826 GIF/ ClFJVsF, IR 813 - 830 I 0 + C1F2 ClF„SbF, Raman 810 385 833 / 6 GIF/ ClF0SbF, IR 803 375 830 2. o + BrF2 BrF„SbF Raman 706 366 - / o + BrF2 BrF„SbF, IR - 365 715 / o ) 107 (3) As a "rule of thumb", the bending mode in bent triatomic species always seems to occur at a wavenumber equal to approximately one-half the symmetric stretching frequency (see Table II-7 in reference 59). In accordance with the assignments of the CIF^* stretching frequencies, the assignment of the bending mode at 544 cm 1 would be too high. Principal difficulties are encountered in the assignment of the vibrational frequencies and for BrF2+, because both are expected and found roughly in the same region as and v^CSbF^ ). This leads to poor resolution noticeable particularly in the IR spectrum. The bending mode \)^{BrF^+) is found at 366 cm thus lending support to the reassignment of V2(C1F2+) quite well. The intense vibration at 706 cm 1 in the Raman spectrum of BrF2SbF^ is best assigned as and a strong shoulder in the infrared at 715 cm 1 is assigned as which is not resolved in the Raman spectrum. This tentative assignment is in good agreement with the C1F2+ bands. The extremely high intensity of v^(BrF2+)^ which is found to be the strongest peak in the Raman spectrum is noteworthy. 6.3.3 Assignment of the Anion Vibrations Assignments for the AsF^ anion in both the IR and Raman spectra of ClF2AsF^ are shown in Table 21. A comparison of the AsF^ frequencies in ClF.AsF,. and in KAsF,, ClOJVsF,. and in a number of complexes 2 b 6 2b containing nitrogen fluoro-and oxy-cations is made in Table 22. The polyfluoroanion complexes containing heterocations of nitrogen have been interpreted as ionic complexes where appreciable anion-cation interaction is absent: small splittings in the mode and the appearance of \>2 as a weak absorption band in the IR spectrum are possibly due to site symmetry effects.'''4''' TABLE 21: Frequency Assignments of AsF in ClF„AsF (cm *) 6 2 6 I.R. RAMAN Assignment 693 vlCAlg) 605,560 603,544 v 2CE ) 688 - v3CFlu) 383 - v4(Flu) 382,372 V5(F2g) F V 2J In contrast to the vibrational spectra of the other AsF.. 6 compounds in Table 22, UpfAsF^ ) appears to be widely split at 603 cm * and 544 cm * in the Raman spectrum as well as in the IR spectrum of ClF^sF^.. v in the Raman spectrum appears to be split at TABLE -22: Fundamental Frequenc ies for the AsF, Anion (cm ) 6 Compound Ref. NOAsF, 693 582* 695 385 379,370 94 0 NOJVsF- 693 582* 690 385 380,370 94 z o N2F3AsF6 690 583* 690 380 376,370 94 N_FAsF, 688 583* :715 - 375 99,100 z o NF.AsF 687 581* 709 406 378 98 4 6 ONF„AsF, 689 584* 692 373 378,372 94 z 0 KAsF, 692 580 698 382 375 THIS WORK D ClOJVsF, 685* 572* 725* 382 372* THIS WORK Z D ClFJVsF, 693 603,544* 688 383 382,372 THIS WORK z o * denotes a vibration found both in the IR and-Raman spectrum 110 382 and 372 cm 1, but the 382 cm 1 absorption may solely belong to the bending mode of CIF^*- The weak absorption band in the infrared spectrum at 310 cm 1 is assigned to the combination band - v^. The band at 520 cm 1 as well as bands at 1290 and 1045 cm 1 in both the IR and Raman spectra are due to impurities of ClO^AsF^ . Weak bands which can be attributed to ClO^* impurities are found in both the spectra of ClO^AsF^ and ClO^SbF^, and presumably arise from hydrolysis or glass interactions. Assignment of the vibrational frequencies for the SbF^' group in ClF_SbF. and BrF„SbF, is more difficult due to the complexity 2 6 2 6 r J of the spectra. In both the I.R. and Raman spectra, far more lines are observed than expected for an anion of octahedral symmetry. Before attempting an analysis of the spectra, it is useful to examine the known crystal structure of BrF2SbF^ reported by Edwards 137 and Jones . Some of their results are reproduced in Figure 22. The structure of BrF^SbF^ consists of polymeric coiled chains with 2 molecules in the chain per unit cell. Each SbF,. unit is distorted o by cis-fluorine bridging to two bromine atoms resulting in the lowering of the anion symmetry from 0^ to At the same time, two SbF^ units are coupled by a single BrF2 unit resulting in a distorted square planar coordination for bromine. While the difference in the Br-F terminal and bridging distances is large (0.60 A), the difference in the same for Sb-F is small (0.07 A), suggesting that BrF^SbF^ remains essentially an ionic structure during the presence of fluorine bridging. FIGURE 22: Crystal Structure of B Fz Sb!| Bond angles (°) R4)-Sb-F(41) = 93 F(3)-Sb—F(3') = 174 F — Br — F = 93.5 o Interatomic distances (A) Sb- F(1) =1.91 Sb— F(4) =1.83 Br— F(1 ) = 2.29 Br— F(2) = 1.69 A.J. Edwards and G.R. Oones G. Chem. Soc. (A) 1467 (1969) 112 When Lta^SbF^. is dissolved in BrF^, these bridges are broken and discrete ions are produced. It is interesting to note the analogy between LM^SbF^ and the white chloryl compounds which when dissolved in HSO^F produced red chloryl cations. The salient features of the crystal structure which may be used in the interpretation of the vibrational spectra of BrF^SbF^ are correlated in Table 23. The lowering of the symmetry of the SbF,. 6 anion from 0^ to C as a result of cis-fluorine bridging results in the splitting of all degenerate modes. In addition, the arrangement of BrF2SbF^ in polymeric coils with two molecules in the same chain per unit cell leads to the overall symmetry D^. The effect of this coupling where the BrF group acts as a spring between two SbF 2 . 6 groups results in twice the number of bands. Assignments for the anion frequencies in both BrF2SbF^ and ClF2SbF^ are made in Table 24, where the vibrational frequency assignments for SbF^ in KSbF^ and ClG^SbF^ are included as a means of comparison. In the case of BrF2SbF^, the bands originating from the mode can be explained as follows: The mode is split into four components, namely, A, B^, B2 and B^. The B modes, B^, and B^ are observed in the infrared spectrum at 535, 523 and 495 cm 1 respectively, while the A mode is infrared inactive. Only three of the four expected modes appear in the Raman spectrum, 552 (A), 523(B2) and 493 (B-r) cm ^ > and it is assumed that the B band in the Raman is 3 1 obscured by the very intense band at 552 cm 113 TABLE 23 Correlation diagram for the XF, species Free molecule Distorted molecule Molecule in chain Point group 0 C„ D • h 2v 2 Vibrational mode: Vl Alg Al A + B2 A A + B v„ E A 2 g B2 B3 + Bl A A + B V3 (°r V Flu Bl B3 + Bl B2 B3 + Bl • A A + B 5 2g A2 A + B2 Bl B3 + Bl TABLE 24 Fundamental frequencies for the SbFA anion. 5 Compound CM -SbF, BrF„SbF, C£02SbF6 KSbF, 2 o 2 6 o Method IR Raman IR Raman IR Raman IR Raman Assignment: [ 660 678 680 680 661 vi V 552 592 535 582 596 586 588 523 523 575 v2 IEg] 542i 495 493 537* 677 673 641 720 706 655 647 644 663 670 649 604 638 640 270 265 267 265 270 269 290 292 285 284 305 308 294 283 282 270 270 288 288 278 The remaining A modes originating from and can be assigned to the strong bands at 678 cm-1 and 641 cm-1, in the infrared spectrum appears to be split into three components at 673, 663 and 638 cm Lack of resolution, however, prevents the observation of the coupling effect for and also for and in which the latter is split into two components at 285 and 270 cm It is very difficult to assign the remaining B modes due to + the closely spaced A modes and v^(BrF2 ) at 706 cm *. Only a weak shoulder is found at 690 cm ^ which is tentatively assigned as originating from v^. The above mentioned high intensity of the 706 cm ^ band suggests a coupling between all the different A modes. It should also be mentioned that coupling will likewise affect + the BrF2 vibrations and for the chain modes, four components for v1 and Vj are expected of the type A, B2 and B . The lack of resolution in the infrared spectrum and the closely spaced bands in the Raman spectrum do not allow a complete assignment beyond the point where the 706 cm 1 band is assigned as the A mode and the 715 cm" shoulder as B2 or B^. It may also be possible that the shoulder at + 690 cm * is better assigned as B^ and belonging to BrF2 . The vibrational spectra of ClF^bF^ are found to be relatively simple. Here, v2(E ) is split into two components A^ and B . 1 1 The B2 mode is observed at 596 cm in the Raman and at 582 cm in the infrared spectrum. The splitting of the Raman active A^ at 542 and 537 cm ^ is interpreted as caused by the crystalline field. Three components of are found at 604, 647 and 677 cm ^ in the IR, the band at 647 cm ^ being also weakly Raman active. v^(SbF^) appears not to be split and is found at 662 cm ^. .116 The above mentioned splittings suggest symmetry for the SbF^ group indicating fluorine bridge bonding over cis fluorine atoms in the SbF^ unit as found for BrF^SbF^. The apparent dissimilarity between the vibrational spectra for both compounds may be due to different crystal structures, or weaker bridge bonding, preventing the coupling of vibrations. 6.3.4 Conclusions and Summary. It is apparent that the anion in the complexes ClF^AsF^, ClF^SbF^ and BrF^SbF^ has less than octahedral symmetry. The distortion of the anion octahedron, more obvious for hexafluoroantimonates than for hexafluoroarsenates, is best explained as due to anion-cation interaction which occurs via fluorine bridges. In general, there seems to be substantial interaction between the anion and cation for _AsF^ and -SbF^ complexes of heterocations of group 7A elements. Covalent interaction has also been postulated 142 by Peacock and co-workers for -AsF^ and -SbF^ complexes containing 143 di-fluoro-phosphorus heterocations and by Gillespie and co-workers for complexes with heterocations of sulphur. In contrast to these heterocations of main group elements in the 3rd and 4th periods (Table 18) the heterocations of nitrogen in the 2nd period are reported to exist as discrete ions in -AsF^ and -SbF^ complexes.1^ 117 CHAPTER SEVEN BROMYL COMPOUNDS 7.1 INTRODUCTION The difficulty in preparing compounds containing bromine-oxygen bonds is well known. The chemistry of the reported bromine oxides, 144-146 146-150 , . . . . 151 ,D . , Br20 , Br02 , and polymers of formulae , (Br20r.Jn and (Br^Og)^ is virtually unstudied. All are of very low thermal stability. Until recently, the only stable species containing solely bromine and oxygen had been the bromate anion, BrO^ . In the past, attempts to prepare the perbromate anion have 152 failed. The non-existence of BrO^ was explained by Urch as due to the absence of good TT bonding between the 2p orbitals of oxygen and 4d orbitals of bromine. In striking contrast to theoretical predictions, however, Appelmann ^3 has recently prepared perbromic acid, HBrO^, in solution and potassium perbromate, KBrO^, a solid stable up to 280°C. Perbromyl fluoride, BrO^F, has also been synthesized1^' and, on the basis of vibrational spectra, appears to have a molecular structure similar to ClO^F with C symmetry. This work reports the attempts to prepare bromyl fluorosulphate, + BrO^O^F, and to investigate the existence of the bromyl cation, Br02 . 1 9 1 The only known Br02~ derivatives FBrO^^f' *' and BrO^O,, ^ are of very limited thermal stability, being comparable to the bromine oxides. 7.2 EXPERIMENTAL An excess of S„0,F. was distilled in vacuo on to 6.503 mmoles z 6 z of KBrOg- Upon warming to 0°C, the reaction proceeded moderately fast with the formation of a bright orange intermediate. The reaction mixture was kept at 0°C. until oxygen evolution has ceased. After removal of the oxygen and the excess ^20^^ by pumping, 6.50 mmoles of KBr(0S02F)^ were obtained. A decomposition point of 170°C. was found. The sample was analysed for bromine after hydrolysis. A second run performed at room temperature led to impurities of Br(S03F)3 and KS03F in the main product which was again KBr(OS02F)4. These impurities presumably arose from the slight decomposition of KBr(OS02F)^ due to the exothermicity of the reaction. Br(S03F)3 was removed from the product by extraction with 820^2 and identified by analysis; KS03F was detected in a Raman spectrum of the product. A third run was made at -40°C. over a period of several hours. The removal of $20^2 at -40°C. left an orange-red solid. On warming to higher temperatures, the solid decomposed with the liberation of O2 and Br2• The residual solid, however, was found to contain large amounts of unreacted KBrOy The absence of $20^2 in the residual peroxide collected from 19 all reaction runs was indicated by F.';NMR and infrared spectra. The residual S„0^F„ was often coloured orange due to very small z 6 z amounts of BrS03F. This orange colour disappeared on standing as BrS03F was oxidized to Br(S03F)3- Attempts to study the KBrO^ - ^O^F,^ """n ^•LUorosu-LPnuric acid were unsuccessful, as KBrO^ alone reacts rather violently with HSO^F to produce KSO^F, Br2 and 0^. 7.3 DISCUSSION The reactions of So0,Fo with KC107 and KI0„ (chapter 9) result Z O Z J 6 in the formation of ClC^SO^F and IO^SO^F respectively. By analogy to these reactions, it was hoped to prepare BrO^SO^F through the reaction of S^O^F'2 with KBrO^. However, this reaction proceeds to the formation of KBr(OSO^F)^. The net reaction is best described as: KBr03 + 2S206F2 *• KBr(0S02F)4 + 3/2.02 It seems feasible that the observed intermediate is BrC^SO^F, expected to be thermally unstable in accordance with the instability of known bromyl compounds. A possible reaction sequence in agreement with experimental observations is: KBrO„ + So0,F„ * KSO-F + Br0_S0„F + 1/2 0_ ozoz J Z J Z Orange Br02S03F * 1/2 Br2 + 02 + 1/2 1/2 Br„ + 1/2 So0,F_ • BrS0_F Z Z o Z o BrS03F + S206F2 • Br(S03F)3 Br(S03F)3 + KS03F »• KBrCS03F)4 120 161 It is interesting to note that the reaction of with KBrO^ is analogous to the KBrO^ - S^O^F^ system, and at room temperature proceeds to the formation of KBrF^. Also, the preparation of ClO^F from KCIO^ and BrF^ is not analogous to the reaction of BrF^ and 162 KBrO^; in lieu of Br02F, KBrF^ is produced along with Br^ and 0^ In exploring the possibility of preparing Br02S03F, it is useful to examine the preparatory routes to Br02F. The first successful/ preparation of Br02F 1^,159 invoivecj the reaction of F2 with Br02 at -50°C.: C5F12 Br09 + 1/2 F + BrO F 1 1 -50°C. 1 Subsequent preparations invoked the use of BrF,. as a fluorinating agent: -61°C. KBr03 + BrF5 y KBrF^ + Br02F + 1/2 02 and BrF5 + 2Br2 + 10 03 + 5Br02F + 10 02 A possible route to the synthesis of Br02S03F might be the reaction of Br02 and S206F2 at -40°C. In view of the results of the study of the KBr03 - S206F2 reaction, it is highly possible that Br02S03F is unstable above temperatures of -40°C. Structural studies seeking + the existence of the Br02 cation and employing the techniques used in Chapter 3 (conductivity studies etc.) would be very difficult indeed;. 121 158 159 It is reported that bromyl fluorocomplexes are not obtained ' from the reaction of Br02F and BF^, AsF,. or SbF,.. This precludes + the possibility of studying the Br02 cation in polyfluoroanion complexes. While the KBrO„ - S„0,F. reaction is unsuccessful insofar as it o Z o Z does not lead to bromyl fluorosulphate as the final product, it does provide a convenient route to the synthesis of K[Br(OSC^F)^]. • The previously reported preparation involving the reaction of KBr and 75 S20^F2 requires a reaction time of several days and heating to 50°C. As no structural information was available on the complex K[Br(OSC^F)^], it became interesting to obtain its Raman spectrum and the spectra of a number of related compounds. The results of these studies are described in the next chapter. 122 CHAPTER EIGHT RAMAN STUDIES OF Br(S03F)3, ICSC>3F)3 AND THE ANIONS [Br(S03F)4]~ AND {USO^]' 8.1 INTRODUCTION The reaction of bromine and iodine with excess S~0,F„ has been / D 2 found to yield the tris-fluorosulphate compounds Br(SC>3F)3 and 3 J3 163 I(SO F), respectivel75 y . Under similar conditions, KBr and KI have been found to give complex salts of the composition K[Br(S03F)4] 43 1 4> and K[I(S03F)4] respectively. Subsequent studies ' ^ ''"^ with the tris-fluorosulphates have yielded some information on their chemical and physical properties as well as their solution .behavior in HSOjF. However, no structural information on any of these compounds has yet been reported. A convenient route to the synthesis of K[Br (S03F)4] was found and described in the previous chapter. It is the object of this chapter to investigate the structure of K[Br(S03F)4] by means of laser Raman spectroscopy, and also the structures of the related compound K[I(S03F)4] and the tris-fluorosulphates Br(S03F)3 and I(S03F)3. The high reactivity of these compounds has prevented infrared studies so far. By analogy to the square-planar anions ICl^ and BrF^ 167,168 the complex anions [Br(S03F)4]~ and [I(S03F)4] are likely to have a square-planar configuration. The most feasible structural models for the tris-fluorosulphates are provided by analogues from among the interhalogen compounds of the AB3 type: (a) the T-shaped monomeric molecule as found for BrF^*^ and CIV^®. (b) the fluorosulphate 172 bridged dimer or possibly polymer similar to the I2^1^ structure. 19 The fact that the F NMR spectrum of the melt or the solution in ^2^6^2^^ shows only a single line for both tris-f luorosulphates might imply that neither model is satisfied, but rapid intermolecular or intramolecular exchange is a definite possibility. 8.2 EXPERIMENTAL The complex K[Br(S03F)4J was prepared from the reaction of KBr03 and $2^6^2 as described in the last; The reaction was performed at 0°C. to avoid impurities of KSOjF and Br(S03F)3 in the product. The sample of K[I(S03F)4] was obtained from Mr. S.P.L. Jones, whose assistance in this project, is gratefully acknowledged. Br(S0,F)„ was obtained from the reaction of BrS0„F and S_0,F„ •5 6 o Z 6 Z 163 and IfS0_F)„, from the reaction of I„ with S„0,F„ . Sublimed o o Z Z o Z analytical grade was supplied by B.D.H. KBrF4 was prepared by the reaction of KC103 and BrF3 as according 3 to Woolf , as a byproduct in the synthesis of ClO^F. 124 8.3 DISCUSSION 8.3.1 Results The Raman spectra of K[Br(S03F)4J and K[I(S03F)4J are listed in Table 25, together with the estimated intensities and tentative assignments. The spectra for Br(S03F)3 and I(S03F)3 are listed in Table 26. Approximate polarisation ratios for the most intense peaks of iodine tris-fluorosulphate were also obtained since the spectrum was obtained for the supercooled melt. All four compounds were found to give well resolved spectra at rather low sensitivities. The Raman spectra of K[I(S03F)4J, K[Br(S03F)4J and Br(S03F)3 are shown in Figure 23. All attempts to obtain well resolved infrared spectra by avoiding window attack did not lead to satisfactory results. 8.3.2 K[Br(S03F)4] and K[I(S03F)4J The Raman spectra of K[Br(S03F)4J and K{ICS03F)4J are very similar, indicating structurally related compounds. The fact that only three vibrations are found in the sulphur-oxygen stretching frequency region of 1500 cm 1 to 900 cm 1 suggest that all four fluorosulphate groups are identical. Some small splittings or :i the presence of shoulders could be due to factor group splitting. Also, a single band attributed to the S-F stretching vibration is found at about 835 cm 1. TABLE 25: Vibrational Frequencies for the [Hal(S03F)4] Anion KBr(S03F)4 KI(S03F)4 Assignment 1424 mw 1409 m vy A' 1407 mw 1237 s 1250 s v1 A' 1220 w,sh 1222 w,sh 970 m,s 1002 m v4 A' 834 m 837 m \>_ A' 615 vs 620 vs v3 A* 578 ms 582 m v5 A' 553 w 554 m vg A" 447 s 442 s VjHal-0 A 406 w 407 w vQ A" 399 w 397 w \>4Hal-0 B2g 270 s 260s \>2Hal-0 B 239 vs 239 vs vfe A' 126 TABLE 26: Vibrational Frequencies for Br(S03F)3 and I(S03F)3 I(s Assignment Br (SO^F) 3 JZ •"•""3 ^ Terminal Bridging 1490 mw 1469 m v? 1467 w 1 1372 m 1381 mw v, '1 1356 mw,sh 1241 s 1233 s,p(0.34) v 1230 m,sh 1168 mw 1182 w,sh -7' 1122 m 1076 w,sh 1050 m,p(0.62) v4' 1010 m 963 m,p(0.29) v. 1015 m 859 m 869 p(0.40) v2' 827 vw 826 w 801 m 800 m v2 721 ms 700 w v5* 645 vs,br 642 vs,p(0.49) v3 612 vs 619 p(0.48) v 3 583 ms 580 m,sh vc 563 m 556 m Vg' 551 vw 540 w 540 m,sh vQ ./continued 127 TABLE 26 ... continued I S0 F Br(S03F)3 t 3 )3 Assignment Terminal Bridging 455 s 457 s,p(0.39) v1Hal-0 430 m 430 w,sh V9 408 mw 412 m V9 384 mw,sh 386 mw,sh v4Hal-0 303 vs 290 vs,p(0.46) V6 276 vs 270 s,sh v2Hal-0 225 ms V6 206 m 181 m 148 m p = polarized, sh = shoulder, w = weak, m = medium, s = strong, br = broad vs = very strong, vw = very weak K[Br(S03F)J i 1 1 1 1 1 1 1 1 1 1 1 1 1 1 i 1600 1400 1200 100Q 800 600 400 200 WAVENUMBER [cm-'] FIGURE 23 129 The observed frequencies differ markedly in total number and - 173 85 position of the absorption bands from that of a SO^F ion ' Covalent interaction will increase the number of vibrational modes from 6 for the SO^F ion with symmetry to 9 for the covalent monodentate or bidentate S0_F group, both with C symmetry by the removal of degeneracy for the 3 E modes. All vibrations for C^v and Cs symmetry will be IR and Raman active. The change in band positions is illustrated nicely for the 174 175 monodentate covalent compounds by ^2^s^2 an<^ related compounds ' , where bands at about 1500 and 1250 cm 1 are found, indicating a larger degree of multiple bonding for the remaining terminal sulphur-oxygen bonds. A correlation diagram, where a tentative assignment for the [HalCSO^F)^] ion is made is shown in Table 27. The degree of splitting of the E..modes, best recognized for the S-0 stretching modes, will indicate the degree of covalency in the halogen-oxygen bond. This is well illustrated when the [BrCSO^F)^] anion is compared to the [ICSO^F)^] anion where a slightly smaller splitting for the latter indicates a slightly more polar bond. The vibrational frequencies of the anions can be interpreted as being due to a covalently bonded monodentate SO^F group. The remaining bands at about 445, 395 and 260 cm 1 are most probably due to halogen- oxygen vibrations and for the bromine species, occur at slightly higher frequencies than the iodine species, as expected from the difference in masses. TABLE 27: CORRELATION DIAGRAM OF THE S03F GROUP S03F" ION (C3v Symmetry): Vibrational W W W V4(E) V5^ VE) modes V V V S03sym SF S03sym S03asym SO^asym rock Example: n group: VlCA«) v^A') v^A') v^A') vy(A") v5(A*) vg(A") v^A') vg(A ) Example: [Br(0S02F)4r (cm_1) 1237,1220 834 615 970 1424 578 553 239 406 1407 0* o 131 The number of vibrational modes agrees well with a square planar _ 167 168 configuration, already suggested by the structure of the BrF^ anion ' which is found to have symmetry. To obtain some indication of the approximate position of the vibrational bands in the fluorosulphate compounds, the Raman spectrum of KBrF^ was recorded which shows a pattern of three absorption lines at 532 (s), 457 (ms) and 246 (m) cm ^. Due to the thermal instability of bromine-oxygen bonds, no other comparable assignments could be found in the literature. The same lack of comparable systems is found for the [IfSO^F)^] anion. All vibrational studies on tervalent iodine-oxygen compounds 176 are concerned with infrared spectra of iodosyl derivatives where the iodine-oxygen bond appears to have appreciable multiple bond character. 8.3.3 Br(S03F)3 and I(S03F)3 The Raman spectra of both tris-fluorosulphates indicate structural similarity but are far more complex than the anionic species. The vibrational frequencies are listed in Table 26, together with a tentative assignment. For both compounds, a-.total number of 6 vibrational modes is found in the sulphur-oxygen stretching region and two strong bands become assignable to sulphur-fluorine stretching. A similar duplication of vibrations is found in the lower range of the bending and rocking modes. Bands previously assigned to halogen- oxygen are found again with comparable intensity in the same place as the anionic species. 132 The obvious explanation for the multitude of bands is that two types of fluorosulphate groups are present in both compounds. One set of bands is in approximately the same place as found for the tetra- fluorosulphatohalate (III) anions, even though the asymmetric S-0 stretching frequency appears at a higher wavenumber. Again, the shift is larger for the bromine compound, indicating a higher degree of covalency as found before. The second set of 9 vibrational modes could possibly be due to a second covalently bonded monodentate SO^F groups as expected for a T-shaped molecule or due to a bidentate bridging group. The latter alternative is chosen for the following reasons: (1) The halogen-oxygen frequencies would be found in a different place and a total number of six Raman active bands would be expected for the C2 model with three in the stretching range. (2) Only minor positional changes for a second terminal SO^F group would be expected. (3) Finally, fair agreement is found with the bridging fluorosulphate 177 group in the hexacoordinated compounds SntSO^F)^ and C^SnfSO^F^ It seems then that the two halogen trisfluorosulphates are not monomeric in the solid state. The most likely models for the structure are a chain-type polymer or a bridged dimer, both with a square planar configuration for the halogen. Even though the observed solubility in solvents such as S20^F^ and ^2^6^2 P0^nts to l°w molecular weight species, no clear distinction can be made at this point. The occurrence of many doublets in the spectrum of BrCSO^F)^ is best explained by assuming different conformations of both the terminal and the bridging group or possibly by factor group splitting. Only singlets are observed for most corresponding bands of I.'GSOjF) present as a melt. All assignments for the two fluorosulphate groups are tentative particularly in the lower frequency range. The higher frequency S-F vibration is assigned to the bridging SO^F group, 177 -1 based on the C^SnCSO^F^ example. Vibrational modes at 200 cm and lower could possibly be due to lattice modes. 134 CHAPTER NINE IODYL COMPOUNDS 9.1 INTRODUCTION A number of inorganic compounds containing the IO2- group have been reported over the years. They include iodyl fluoride, 10 .^F^-^'^, and its reaction products with the Lewis acids, BF^ and AsF^'''"'", where + - 11 199 the latter compound is commonly regarded as IO2 AsF^ ' and not + aS su este 2 AsF4 I02F2 SS d previously ^. The other known iodyl compounds can be regarded as salts of strong oxyacids and include the compounds 201 202 1R9 707, 204 (IO^^O^ , I02S03F , I02CF3C02 , (I02)2Se04 and (IO^HSeO 4 Iodyl compounds are white or pale yellow solids, which were + previously thought to contain the iodonium cation, I02 . The existence 205 of this cation has been postulated as a reaction intermediate + However, no evidence has been found for the existence of I02 in the 206 sulphuric acid solvent system . In addition, infrared vibrational 207 spectra of the related iodosyl salts (10) ,^0^ and (IO)2Se04 have suggested the structures of these compounds to consist of polymerized 1-0 chains cross-linked by the SO^ or SeO^ groups present as either discrete anions or four-covalent groupings. As a consequence, similar formulationr i s for r th+1,e iody• A li saltn-s havv. e resulte-n-d ,203,204,20 6 , bu, t+ any structural proof for the postulated iodine-oxygen bridged structures is clearly lacking. Except for an incomplete infrared spectrum of 199 208 I02AsF^ , and a very brief report on the infrared spectrum of I0,,F , no detailed vibrational spectra of the iodyl compounds has been reported. 135 It is therefore the purpose of this chapter to apply infrared and laser Raman spectroscopy to a systematic study of the iodyl compounds, I02F, I02AsF6 and I02S03F. It should be possible to decide, whether a common structure of the iodyl group is found in all these compounds and whether such a group is polymerized via discrete I-O-I bridges or exists as the cation with multiple bonding from oxygen to iodine, + analogous to the formally related chloronium cation,C102 . Anion-cation interaction, as found in the chloryl compounds discussed in Chapter 4, is also a distinct possibility. Inclusion of the related compounds lOF^ and KI02F2 in this vibrational study should help in the interpretation since the X-ray 209 210 crystal structures are reported for both compounds ' . No detailed vibrational spectra appear to have been reported for these compounds. Solution studies in fluorosulphuric acid were undertaken for three main reasons: 206 (1) The study by Gillespie and Senior uses iodic acid as the sole + solute in I-^SO^. A possible I02 cation, formed according to the net + + equation;:HI03 + 2H2S04 = I02 + H30 + 2HS04~, would have to coexist + with the H30 ion. Solution studies of solutes already containing + the iodyl group would avoid the formation of the H30 ion and would + present a better chance to obtain evidence for the I02 cation. It is interesting to note that, in analogy, a solvated chloronium cation is formed not by solvolysis of metal chlorates ar chloric acid, but by chloryl compounds as solutes (chapters3,4). (2) Fluorosulphuric acid is a stronger protonic acid than H^SO^ and is clearly better suited for the study of such highly electrophilic 27 cations as shown by many examples (3) The most ideal solute in HSO^F, iodyl fluorosulphate, was first 189 reported to be insoluble in HSO^F . It was found later during the course of this research that the compound dissolves at room temperature in HSO^F in an extremely slow process over several days to give clear solutions, concentrated enough for conductivity studies. Finally, the reactions of S^^F,, with KIO^ and KIO^ were investigate to see whether compounds containing the IC^- grouP or ^®_<~ grouP would result, or if the formation of KIG^CSO^F^, analogous to KI02F2, would occur. 9.2.1 EXPERIMENTAL 9.2.1 The Preparation of I02SC>3F Iodyl fluorosulphate was prepared from the reaction of l20j- and 189 ^2^6F2 as accordi-ng t0 Aubke et al 9.2.2 The Reaction between KI0„ and S_0,F„ o z o z An excess of S.CLF,, was distilled in vacuo on to 9.67 mmoles of z 6 z finely powdered KIG"3 (BDH) contained in a reaction vessel with a stirring bar (Figure 1) cooled at liquid nitrogen temperature. On warming to 25°C, a very slow reaction occurred with the liberation 137 of 0^. The rate of reaction was enhanced by raising the temperature to 50°C. - higher temperatures resulted in decomposition of the product and after about 12 hours, the evolution of 0^ had ceased and the reaction was complete. Removal of the excess -^O^F2 ^-e^ Dehind a flaky pale yellow solid of weight corresponding to a mixture of 9.65 mmoles of both IC^SO^F and KSO^F. The presence of these compounds in lieu of KIC^(SO^F^ was revealed by a Raman spectrum and x-ray 19 powder pattern of the product. Gaseous IR and F NMR spectra of the residual peroxide indicated the absence of S20,-F2. 9.2.3 The Reaction between KIO, and S„0,Fn 4 2 o 2 An excess of So0,F_ was reacted with 3.01 mmoles KIO. (BDH) Z o Z 4 under the same reaction conditions as in section 9.2.2. Again, the reaction proceeded very slowly at room temperature with the evolution of 0^, while temperatures greater than 60°C. resulted in the decomposition of the product. To complete the reaction, heating at 50°C. for 12 hours was required. Removal of the excess peroxide left behind a flaky pale yellow solid resembling the product obtained in the KIO^ - ^2*^6^2 reacti°n- A Raman spectrum indicated the product to be a mixture of IO^SO^F and KSO^F, while the weight change corresponded to 2.94 mmoles for each component. In each run, the weight change was found to be less than the theoretical value due to the loss of small 19 amounts of IO2SO2F to the vacuum line by sublimation. F NMR and snowe gaseous IR spectra of the residual 520^p2 d that 820,^2 was not formed in this reaction. 138 9.2.4 The Preparation of KIO F2 Potassium difluoroiodate was prepared by the evaporation of a solution of KIO^ in 52% HF (Fisher Scientific Co.) as according to Helmholz and Rogers21^. 9.2.5 The Preparation of I0F3 g IOF^ was prepared by the reaction of ^^5 anc* boiling IF5 As the description of the procedure reported in the.literature is rather g brief , a detailed account of the preparation now follows: Impurities of iodine and complexes of I^+ in IF^ (giving rise to a blue-coloured liquid) were removed by distilling IF,, from the main cylinder (Matheson Co.) into a storage trap containing about 10 g. of mercury. A metal vacuum line was used in this transfer. After about 50 g. of IFj- had collected, the storage trap was removed from the line and shaken for several minutes to ensure complete reduction of the iodine and species to I in the form of the nonvolatile solid Hgl2- The storage trap was once again attached to the metal line. The glass apparatus used for the preparation of I0F^ is illustrated rev us in Figure 24. About 1.5 g. of finely powdered I20,- > P i° ly heated for 6 hours at 100°C. in vacuo, was admitted to the apparatus 'at the opening B. The vessel was flame sealed at C and attached to the metal line at A via a flexible coil of copper tubing, to allow movement of the apparatus while attached to the metal line. About 30 g. of the 139 B Kontes Valve A * Coarse Filter v E / D FIGURE 24: Apparatus used for the preparation of IOF^ purified IF^ was distilled from the mercury storage trap on to the I^O,. contained in the D compartment of the evacuated reaction vessel. Dry air was then admitted into the system via a ?2®5 drying tube attached to the vacuum manifold. The entire reaction vessel was then immersed in an oil bath and heated until the IF,, began to boil. After several minutes and constant shaking of the vessel, a large amount of ^2^5 nad dissolved in the colourless liquid. The hot solution was then pumped through the coarse filter into compartment E, while the vessel was tipped sideways to aid in the transfer. White needle• like crystals separated from the mother liquor on cooling. The final product was obtained by pumping off the excess IF,, at room temperature. 9.2.6 The Preparation of I02F Iodyl fluoride, IC^F, was prepared by the method of Aynsley g et al. by heating I0F_ at 110°C. in the presence of dry air as according to: 1.10? I02F + IF5 + 140 The white solid was heated at 110°C. under vacuum for a short period of time to eliminate all traces of IF^. All attempts to prepare iodyl fluoride by the direct fluorination 19 of a solution of in anhydrous HF were unsuccessful. A typical experiment is described as follows: A quantity of 6.0 g. of finely powdered and purified 1was dissolved in 20 g. of anhydrous HF (Matheson Co.) contained in a Kel-F trap. Fluorine gas (undiluted) was bubbled vigorously through the solution for a period of 27 hours at 25°C. (Moisture is prevented from entering the trap by a large pressure of fluorine which also reacts with ^0 to convert it to HF.) The flow of fluorine was then halted and the Kel-F tube immediately attached to the vacuum line. The solvent was removed by pumping at 50°C. for several hours. A Raman spectrum of the product showed it to contain a mixture of HIO^ and ^2^5 ON^ * wnile an analysis revealed only trace amounts of fluorine. As this experiment was attempted several times under different reaction conditions and led to either HlOg or unreacted ^^c; in eacn instance, the original literature report must be regarded as highly questionable. 9.2.7 The Preparation of I02AsF6 Iodyl hexafluoroarsenate was prepared by the reaction of AsF^ and a solution of IC^F in anhydrous HF.^ As the literature reports of this preparation are rather brief, the synthesis of IO^AST^ is now described in detail: 141 About 15 g. of anhydrous HF was distilled through a metal vacuum line into a pre-weighed monel metal reaction vessel (Figure 2) containing 1.2 g. of I02F cooled at -196°C, followed successively by about 5 g. AsF,.. The reaction vessel was warmed to room temperature and then heated at 40°C. for two hours. The anhydrous HF and excess AsF^ were then removed by pumping. The resulting product was a yellow powder and was identified as IG^AsFg by analysis for iodine and the weight change. A similar attempt to obtain IO^AsF^ by the reaction of IC^F and AsFr without the use of HF as solvent resulted in a mixture of IO^AsF. o 2 o and unreacted 9.3 RESULTS AND DISCUSSION 9.3.1 The Reactions of S„CLF„ with KIO, and KIO, 2 o 2 3 4 The reaction of KIO^ and S_®(^2 Proceec*s siowly at room temperature to a mixture of KSO^F and IO^SO^F as established by a Raman spectrum of the products. The reaction, analogous to the KC10„ - S_0,F„ 6 2 o 2 reaction, can be formulated as: KI03 + S2°6F2 ' KS03F + I02S03F + 1/2 °2 The reaction of KIO. with S„0,F_ results in the same products , and 4 2 o 2 proceeds as follows: KI04 + S2°6F2 " KS03F + I02S03F + °2 No evidence for the existence of the complex KI02(S02F)2 could be found. 142 9.3.2 Vibrational Spectra of Iodyl Compounds 9.3.2.1 Results Well resolved Raman spectra for the iodine-oxygen compounds were obtained in all cases, except for IC^AsF^, where strong scattering of the laser light occurred. Due to the extreme reactivity of lOF^ and IC^SO^F, satisfactory IR spectra for these compounds could not be obtained. The vibrational frequencies for lOF^ and 0(^2 are listed in Table 28, together with estimated intensities and suggested assignments, and the Raman spectra of lOF^ and 1(1(^2 are reproduced in Figure 25. The vibrational frequencies for IC^F and IC^AsF^ together with crude assignments and estimated intensities are contained in Table 29. The Raman spectra of IC^F and IC^SO^F are shown in Figures 26 and 27, while the infrared spectra of IC^F and IC^AsFg are shown in Figure 28. 9.3.2.2 I0F3 and I02F2~ 209 - 210 The structures of lOF^ and the 102^2 are known and therefore their vibrational spectra should be helpful in interpreting the spectra of the iodyl compounds. Both lOF^ and the 102^2 ~^on ^n ^®2^2 are f°unc* to nave related structures, derived from a trigonal bipyramidal configuration, with a stereochemically active lone pair on the central atom, two fluorine atoms in axial positions and the remaining two atoms in equatorial positions. As is generally observed for this bond type, the equatorial 143 TABLE 28: Vibrational Spectra of I0F3 and KI02F I0F3 KI02F2 Observed Frequencies (cm Observed Frequencies (cm ) Raman Assignment Raman I.R Assignment 918 vw—. 855 V I02 asym 907 w Iodine- 830 w,sh 844 " ] oxygen 883 vs 808 vs 816 mVS -»J stretch I0o sym 862 w 805 vs J z 809 w ^ I02F 472 s 476 s VIF 716 w J Impurity 434 s 657 vs 355 ms 358 m,sh VI-F eq. 550 vs 345 m,sh 340 s VI-F ax. 5l0E 515 ms 324 s 322 m,sh VI-F 350 m -•I 220 ms I Deformation ^2 304 ms 190 m,sh 196 ms j Bands rock 295 vw -J vw = very weak, vs = very strong, ms = medium-strong, w = weak, s = strong, m = medium, sh = shoulder FIGURE 25: Raman Spectra of IOF and 10 F Raman Spectra of IOF and IO F i l 1 1 1 1 1 1 1 10OO 800 600 400 200 145 bonds are shorter and stronger than the axial bonds. Both structures are distorted from the trigona^. pyramidal configuration ( as indicated by the bond angles) as a result of the influence of the non-bonding pair of electrons. Also, shorter bonds are found in lOF^ than for 10^2 , although a comparison is difficult due to the large error limits reported for the latter. The fact that unequal iodine- axial fluorine bond distances in I0Fg are observed is indicative of I...F intermolecular interaction. Nine infrared and Raman active vibrations are expected for I0F with C symmetry, while the I0?F ion., with C„ symmetry, is expected to have 9 Raman active and 8 infrared active modes. The described structural type is rather unprecedented in the chemistry of the main group elements, and only for the related ionic species — 181 SeOFg are vibrational spectra reported. A total of 7 vibrational modes are found for lOF^, and these are tentatively assigned in Table 28. It can be argueddthat the remaining deformation modes are found at lower frequencies and are obscured by the strong background from the exciting light in this region. Low frequency vibrations are in fact found in the far infrared spectrum 1 of KI02F2 at 220 and 196 cm" . A tentative assignment of the vibrational frequencies for lOF^ 190 can in part be based on the vibrational frequency assignments for I0F,. , where polarization measurements on the liquid sample and the observation of numerous combination bands in the infrared have allowed more complete 209 assignments. As indicated by the reported structure of I0F^ , a non-bridging 146 iodine-oxygen bond exists in lOF^, and the observed absorption band at 883 cm '''is assigned to the iodine-oxygen stretching vibration (cf. 927 cm * for 1=0 stretching in IOF,.) . Weak absorption bands at 918, 907 and 862 cm * very possibly arise from factor group splitting. Iodine- fluorine stretching modes are observed at 657, 550 and 515 cm The band occurring at the highest frequency, 657 cm *, is assigned to the equatorial fluorine - iodine stretch. The frequencies of the two remaining stretching vibrations involving the axial fluorine atoms are considerably lower, which may indicate some intermolecular interaction. No certain assignment is possible at this point for the bending vibrations. The observed frequencies for the vibrational spectra of KIO2F2 199 are in agreement with the few reported values . Absorption bands due to the iodine-oxygen stretching modes are again split and occur at considerably lower frequencies than the iodine-oxygen stretch in IOFy The iodine-fluorine stretching vibrations in lO^P^ a^s0 occur at lower frequencies, indicating weaker iodine - fluorine bonds. A strong absorption band observed at 434 cm ^ in the IR spectrum, but not found in the Raman spectrum, is possibly a combination band. A simlar spread of element - oxygen and element - fluorine stretching — 181 vibrations is found for the SeOF^ ion in KSeOF^. No definite assign• ments for the bending modes is presentedciby thedauthorsv.•. 9.3.2.3 I02F Discrete IC>2~ groups in iodyl compounds would be expected to absorb in the same region as the iodine - oxygen vibrations in lOF^ 1 and I02F2~- This is further supported by the value of 810 cm for the 147 TABLE 29: Vibrational Spectra of 10 F, 10 AsF and 10 2 I0 F I02AsF6 J 2°S -1 IR (cm ) Raman (cm ) Assignments IR (cm ) Assignment Raman (cm 858 s 866 w I - 0 774 w,sh 221 w 830 s 830 w,sh stretching 718 s 240 s 801 s 807 vs modes 689 s 276 w 3ASF6 716 s 705 vs 314 w 614 m,w 335 w 557 vs 550 vs I - F 534 w,sh 352 w 534 m,sh stretching 394 m,sh 373 w 384 s v.AsF, 398 w 4 6 356 ms 455 m 364 w Deformation 337 ms 376 m,slv 541 w 328 s modes 300 ms 304 mw 350 w,sh 616 w 251 301 m 644 w 671 w 735 s 747 vs 776 m 798 w 814 w 148 FIGURE 26: Raman Spectrum of KI09F? 149 FIGURE 27: Raman Spectrum of IO-SO F o rO CM 150 FIGURE 28: Infrared Spectra of I0„F and I0_AsF/ 2 2 6 I02F 1000 cm-1 250 10£AsFs 384- 1000 cm-1 250 151 frequency of the symmetric stretching mode of the TeC^ molecule, isoelectronic to 10 ^', as determined by a vibrational analysis of the electronic absorption spectrum*^. The existence of discrete K^- groups is indeed found for iodyl fluoride, as evidenced by the vibrational frequencies in the region of 800 cm ^. The infrared spectrum is in good agreement with the 208 previously reported one . The Raman spectrum shows three very strong absorption bands in the stretching region at 807, 705 and 550 cm . The extreme complexity of bands in the iodine-oxygen region - a total of three bands are approximately described as 1-0 stretching vibrations - indicates extensive vibrational coupling. Simple ionic structural models such as I02+F or I02+I02p2 are not consistent with the observed spectrum. The absorption band at 550 cm "'is in the iodine-fluorine stretching region and is comparable to the frequency values found for the iodine-fluorine stretch in IOF^ , where the fluorine is in an axial position. Fluorine bridge bonding, already indicated by the results of the x-ray analysis of IOF^j may also be involved in the structure of In addition, bridge bonding via oxygen is possible, as suggested by the absorption band at 716 cm ^, which is considerably lowered from the expected range of IO2 stretching frequencies. (Frequency values of about 600 cm ^ are 207 assigned for I-0-I vibrations in iodosyl compounds .) Both types of bridge bonding could account for the apparent polymeric nature of iodyl fluoride . Complex vibrational spectra are also found for iodine 192 pentoxide, where in spite of a previous attempt using infrared data , no simple assignments appear possible. The Raman spectrum of I20<- has been recorded in this work, and the frequency values are listed in Table 29. 9.3.2.4 I02AsF( By analogy to the chemical behavior of other oxyfluorides, e.g. ClOJF^^^'^^, the reaction of IC^F with strong Lewis acids such as BFg and AsF^ should proceed under fluoride abstraction and + I02 ion formation. However, the infrared spectrum of IG^AsF^ shows no absorption bands between 800 - 1000 cm 1, the region expected for + 199 I02 stretching frequencies, in agreement with previous work Bands at 689 and 384 cm 1 are in agreement with the expected frequencies for the and modes of the AsF^ ion, but the highest band observed is a weak shoulder at 774 cm ^. Two possible explanations can be offered for the observed spectrum (1) The structure of IO^sF^, consists of polymeric cations with iodine-oxygen-iodine bridge bonds and AsF^ ions. Absorptions at 614 and 535 cm 1 could be attributed to I-0-I bridging. (2) Complex formation between AsF,. and I02F proceeds via oxygen 187 bridging as found recently for SeOF^SbF,. . However, it would be surprising to find As-F stretching and bending modes in the expected positions of and for the AsF^ group. The failure to observe a Raman spectrum for this compound does not allow a more detailed assignment of the listed vibrational modes. It can be concluded that + a monomeric I02 cation seems to be absent. 153 9.3.2.5 I02S03F A different situation is encountered in the Raman spectrum of IO^SO^F as shown in Figure 27. Strong absorption bands are found at 876 and 865 cm * in the region expected for a monomeric iodonium cation. However, the observed spectrum has far too many absorption bands to allow a simple interpretation for the structure as IC^SO^F , even if allowance is made for factor group splitting. The sulphur-oxygen stretching region shows three main bands at about 1335, -1170 and -1030 cm * with extensive splitting for the latter two bands. This is indicative of a symmetry lowering for the SO^F group from C^v in the free ion to C . Such symmetry lowering would be expected when the SO^F group becomes covalently bonded to the cationic group and acts either as a monodentate group or as a bidentate bridging or chelating group. As discussed in Chapter 8, bidentate grouping are found for Br(S03F)3 and I(S03F)3, as well as monodentate S03F groups. The frequencies for the bidentate groupings in BrCS03F)3, ICS03F)3 and 177 49 193 also in (CH3)2Sn(S03F)2 ' where the crystal structure is known , are listed in Table 30 with the observed vibrational frequencies of I02S03F. As can be seen in the table, good agreement exists between the frequencies of the reference compounds and I0,,S03F. It is on this evidence that the characteristic S03F frequencies in the spectrum of I02S03F are assigned to a bidentate bridging grouping, rather than a bidentate chelating grouping, resulting in a polymeric structure. TABLE 30: Vibrational Frequencies of Bridging SO^F Groups '(em *) I0 S0 F Assignment I(S03F)3 Br(S03F)3 (CHj)2Sn(S03F) 2 3 (Raman) (Raman) (IR) (Raman) 1 1381 1372,1356 1360 1336 ,1332 1182 1168 1180 1172 7 1076,1050 1120 1088,1072 1070 4 869 859 825 865 2 700 721 725 692 5 642 645 650 663 3 556 563 555 565 8 430 430 420 431 9 290:: 303 304 318 6 155 Assignments for the IC^ group can now be made. The bands at 900 and 876 cm ^ are assigned to IO2- asymmetric stretching and the band at 865 cm ^ to 102~ symmetric stretching. The occurrence of both the asymmetric and symmetric stretching frequencies of the cationic group in the Raman spectrum indicates a non-linear arrangement for the IO2- grouping. A band at 459 cm * is most likely to be the 10^ bending vibration. The very strong band at 428 cm * is assigned to a symmetric stretching involving the singly bonded oxygen atoms and the central iodine atom. The band at 520 cm ^ is assigned to the corresponding 0- 1-0 asymmetric vibration, again indicating a non-linear arrangement, as found in a distorted trigonal bipyramidal structure with a stereo• chemical ly active lone pair on iodine. In conclusion, a polymeric structure is suggested with cation IO2 groups, where strong iodine-oxygen multiple bonding is present, and polymerization via bridging anions,rather than bridging over 1- O-I groups, occurs. The proposed structure for I02S03F is illustrated in Figure 29. A principal structural difference between IO2SO2F and I02AsF^ is quite apparent. 9.3.3 Solution Studies Iodyl fluorosulphate, originally reported to be insoluble in 189 HSO^F , dissolves slowly in this solvent over a period of several days with a noticeable increase in conductivity. The specific conductivities of I02S03F in HSO^F are shown in Table 31 and plotted _2 versus molality in Figure 30. At concentrations less thanr3.0 x 10 molal, FIGURE 29: The Proposed Structure for I02S03F TABLE 31: Specific Conductivities of I02S03F in HS03F at 25.00°C. KS03F Additions 102 x Molality 104 x K 102 Molality 104 x K (fl 1cm 1) C^_1cm-1) 0.0000 1.125 0.0429 31.45 0.0536 1.912 0.3477 36.45 0.2882 5. 728 1.210 51.86 1.026 13.83 2.796 83.01 1.541 18.07 3.725 101.6 2.071 21.68 4.911 125.5 2.778 25.92 3.196 28.49 3.635 30.77 Interpolated Specific Conductivities 4 -1-1 K x 10 (fl cm ) 7 KS0 F 10 SO F Molality x 10 6 0.00 1.085 1.125. 0. 25 7.0 4.9 0.50 13.61 8.0 0.75 19.7 10.8 1.00 25.8 12.4 1.50 38.0 17.7 2.00 50.0 21.5 2.50 62.5 24.8 3.00 72.7 27.7 3.50 84.9 30.1 4.00 97.0 31.9 120-i o KSQ.F standard Electrical conductivity of IQ, S03F_ 0 IO2SO3F in HSQF at 25°C. A KS03F added 100H E (J 80^ t o x ^ 60 "d c o o g 40 a CO 4CH Molality xio2 moles/kg 159 IG^SO^F behaves as a weak - medium base in HSO^F, presumably giving + rise to solvated I02 cations and SO^F anions. However, with increasing concentration, smaller increases in conductivity are observed , possibly because of the formation of polymeric species. The addition of KSO^F to the solution results in an increase in conductivity, but the slope of the KSO^F addition curve is not quite as steep as expected for the normal reference KSO^F conductivity curve. It is unlikely that IC^SO^F is exhibiting acidic behavior to account for the slightly reduced conductivity values for KSO^F, but it is very possible that fluorosulphate anions are being consumed in polymer formation. It seems that at low concentrations, IC^SO^F can dissolve in + HSOgF and form solvated I0"2 cations by the breaking of bridge bonds. At higher concentration, dimerization followed by higher polymerization probably occurs. However, the nature of these polymeric species was not determined as the solutions were too dilute to obtain meaningful Raman spectra. 160 CHAPTER TEN SELENIUM (IV) OXYFLUOROSULPHATE 10.1 INTRODUCTION This chapter describes the synthesis, properties and reactions of the new compound, selenium (IV) oxyf luorosulphate, SeOfSO^F^. In addition, the structure of SeOfSO^F)^ is discussed on the basis of its Raman vibrational spectrum and conductimetric measurements in HS03F. No other oxyfluorosulphates of selenium have yet been reported, and in fact few fluorosulphates of selenium are known. Barr et al. 178 179 have prepared and characterized the ionic compound Se^fSO^F)^ ' The fluorosulphate, SeF^SO^F, prepared by the addition of SO^ to 21 SeF^, has been known for some time , but structural studies have appeared only recently1^. 10.2 EXPERIMENTAL 10.2.1 Materials Reagent grade selenium oxychloride and selenium dioxide were used without further purification. Analytical grade selenium was finely ground in a mortar and pestle. Potassium oxotrichloroselenate (IV), 181 KSeOCl^, was prepared as according to Paetzold and Aurich (anal %C1 Theor: 44.2, Found: 43.6) 161 Bromine (I) monofluorosulphate was prepared by the method of 42 Roberts and Cady . Anhydrous analytical grade stannic chloride was obtained from J.T. Baker Co. and further purified by distillation. 45 Potassium fluorosulphate was prepared as according to Barr et al. The transfer of all compounds was performed in a dry-box, while a high vacuum system was employed in the preparation of selenium (IV) oxyfluorosulphate. 10.2.2 Preparation of SeO(S03F)2 Three routes leading to the synthesis of SeOfSO^F)^ were found and are described as follows: (1) An excess of S^^F,^ WaS distilled in vacuo into a 150 ml pyrex reaction flask (Figure 1) containing in a typical preparation 28.31 mmoles of SeOCl,, at -196°C. Upon warming the reaction to room temperature, a mildly exothermic reaction took place with the evolution of Cl,,. At regular intervals, the Cl2 was pumped off to avoid the formation of CISO^F and possible side reactions, and the reaction vessel was shaken. To complete the reaction, fresh S2P^F2 was added and the reaction vessel shaken until the evolution of Cl,, had ceased, and the pale yellow colour of SeOCl2 had disappeared. The excess S20^F2 was then removed, leaving a colourless, viscous and non-volatile liquid, corresponding to 28.25 mmoles of Se0(S03F)2> Anal. Calc. for SeO(S03F)2: Se 26.94, F 12.95; Found: Se 26.93, F 12.59 162 (2) An excess of S„0,Fo was distilled in vacuo on to 9.27 mmoles of 2 6 z Se02 contained in the usual type of reaction vessel(Figure 1). Upon warming to room temperature, a slow reaction took place with evolution of 0^. The reaction vessel was heated at 50°C, and after 10 hours had elapsed, all solid Se02 was consumed, and a viscous liquid, insoluble in $20^2, remained. The excess S_®(f2 was pumped off at room temperature, leaving 9.24 mmoles of SeOtSO^F^ behind. The product was identified by its Raman spectrum and Se analysis (26.3% v.s. 26.9% calc). (3) A small excess of BrSO^F was distilled in vacuo on to 6.60 mmoles of Se0Cl2 contained in a pyrex reaction vessel (Figure 1). At room temperature, a moderate reaction proceeded with the evolution of and CI2. Removal of the excess BrSO^F left behind a viscous liquid, contaminated with BrSO^F. The Raman spectrum of the product and the mass change was in agreement with the formation of SeOfSO^F^. Because of the difficulty in removing the excess of BrSO^F from the product, the first two preparations have always been employed in further preparations of SeOfSO^F^. 10.2.3 Attempts to Prepare KSe0(S03F)3 (1) Into a reaction vessel completely sealed off from the atmosphere and containing 3.41 mmoles Se0(S03F)2> milligram quantities of KS03F were added with shaking, until 3.40 mmoles had been added. At room temperature, no visible reaction was observed, and warming 163 the reaction mixture to 60° - 80°C. produced no apparent effect. A Raman spectrum of the mixture only showed the original components. (2) An excess of S„0,F„ was distilled at -196°C on to 1 g of KSeOCl, z o Z • ° 3 contained in a reaction vessel similar to that used in the preparation of Se0(S0.jF)2 (Figure 1). At room temperature, a reaction took place immediately to form a white solid, a viscous liquid and chlorine gas. The mass change was consistent with the formation of a product of the composition of KSeOtSO^F)^; a Raman spectrum of the product, however, could be interpreted as a mixture of KSO^F and SeOfSO^F)^. (3) Highly purified HSO^F was added to 1 g of KSeOCl^ contained in a 2-part pyrex vessel and equipped with a Fischer and Porter teflon valve. A yellow solution immediately formed and efforvescence occurred. A Raman spectrum showed the presence of SeOCl2. Evaporation of the liquid phase left residual solid KSO^F as according to a Raman spectrum. 10.2.4 The Reaction Between SnCl4 and SeO(S03F)2 An excess of SnCl^ was distilled in vacuo on to 1.4 g SeO(S03F)2. After the reaction flask was shaken for several minutes at room temperature, the excess SnCl^ was distilled off, leaving a glassy solid behind. The colourless glass failed to crystallize when it was cooled at -196°C. and allowed to slowly warm up. A Mossbauer spectrum of the glassy product showed a broad peak at 6 = -0.133 mm/sec relative to Sn02. This chemical shift is too 49 low for the product to be a simple adduct of SnCl^ Possibly, a ligand scrambling reaction occurs in which one or more of the 164 chlorine atoms is replaced by fluorosulphate grouping which would account for the low chemical shift. 10.2.5 Vibrational Spectra The Raman cells for liquids and solutions described in chapter two were used. Attempts to obtain an infrared spectrum of SeOCSO^F)^ between AgCl plates resulted mainly in a spectrum of AgSO^F and SeOCl^, evidently as the result of the reaction: 2AgCl + SeO(S03F)2 • 2AgS03F + SeOCl2 The use of silicon windows also failed to produce satisfactory spectra. 10.2.6 Conductivity Studies Conductivity measurements were obtained for solutions of SeO(S03F)2 in both HS03F and H[SbF2(S03F)4]. The preparation of superacid, 2 H[SbF2 (S03F)4J, has been previously described^ ' . 10.3 RESULTS AND DISCUSSION 10.3.1 The Preparation and Properties of Se0(S03F)2 Three synthetic routes have) been found leading to the new compound, SeO(S03F)2. Peroxydisulphury1 difluoride, ^2^6^2' and bromine (I) monofluorosulphate, BrS03F, are employed as fluorosulphating agents as according to: 165 25°C. SeOCl„ + S_0,F. SeO(S03F)2 + Cl2 2 2 6 2 (1) 25°C. SeOCS03F)2 + 2BrCl* (2) 50°C. Se0o + So0,F_ SeO(S03F)2 + 1/2 02 2 2 6 2 (3) 10 hours * BrCl dissociates into Cl„ and Br„ at room temperature The first reaction is certainly the most convenient method of preparing SeO(S03F)2. The difficulty in separating the product from the relatively non-volatile bromine monofluorosulphate prevents the second reaction from being a useful route to Se0(S03F)2, while the third reaction is slow and requires several hours for completion. Selenium (IV) oxyfluorosulphate is a highly reactive compound which reacts violently with water, and explosively with organic material. A viscous, colourless and non-volatile liquid at room temperature, it solidifies to a glass at -45°C. The compound is thermally stable in a pyrex container up to about 60°C. at which point the liquid turns a pale yellow and some decomposition begins. The 3 density of selenium (IV) oxyfluorosulphate is 2.43 g/cm . It is insoluble in CCl^, moderately soluble in S^^F,, and miscible with HS03F. The neat liquid has a temperature-dependent electrical conductivity typical of anoelectrolytic dissociation as is shown in Figure 31 .and in Table 32. TABLE 32 Specific Conductivity of Neat SeO(S03F)2 as a Function of Temperature 3 TEMP.(°C.) K x 10 (fi"1cm_1) 19.15 1.228 25.00 1.592 28.00 1.831 29.00 1.905 33.20 2.235 36.70 2.545 40.40 2.835 5 2U 1 1 cf. SeOCl2 at 25.00° C. K = 2.00 x 10 n cm Conductivity of neat SeO(S03F)2 as a Function of Temperature Temperature ( °C) FIGURE 31 Acid-base properties of the covalent molecule, SeOCl^j are well 184 known . e.g. Basic behavior is exhibited by the donor oxygen atoms in Se0Clo which can interact with the Lewis acid SnCl. to _ 2 4 185 186 form the cis-addition compound SnCl^.2SeOCl2 ' . On the other hand, SeOC^ can behave as a Lewis acid in reacting with alkali metal 181 chlorides to form compounds of the type MSeOCl^ , where M is an alkali metal. In a similar fashion, SeOF^ also forms compounds of 181 the type MSeOF^ with alkali metal fluorides . (However, little is known about the possible donor properties of SeOF2, other than 187 the very recently reported crystal structure of SeOF2-NbF,_ where the oxygen atom of SeOF2 is bridged to niobium.) Selenium (IV) oxyfluorosulphate differs markedly from SeOC^ in donor-acceptpcr^properties. All attempts to prepare KSeO(SO^F)2> the analogue of KSeOCl^, were unsuccessful and it is unlikely that this compound exists as evidenced by the following: (1) The addition of finely ground KSO^F to SeO(S03F)2 does not result in a reaction. (2) The reaction of KSeOCl_ with an excess of S_0,Fo j 2 6 z produces a mixture of SeO(S03F)2 and KSO^F as according to: KSeOCl3 + 3/2 S206F2 • KS03F + SeO(S03F)2 +3/2 Cl2 (3i) A mixture of SeOCl2 and KS03F results from the reaction of KSeOCl3 and HS03F: . KSeOCl3 + HS03F > KS03F + SeOCl2 + HCl 169 Possible donor properties of SeOCSO^F)^ were investigated by reacting it with SnCl^. A glassy solid of uncertain composition resulted instead of a simple addition compound. The Mossbauer spectrum revealed a broad peak at 6 = -0.133 mm/sec relative to Sn02, far outside the normal range of +0.2 to +0.5 mm/sec for adducts of 119 SnCl. . It is feasible that one or more chlorine atoms in SnCl„ 4 4 are replaced by fluorosulphate groups accounting for the negative isomer shift, and that SeOfSO^F^ does not behave as a Lewis base in this reaction, but as a fluorosulphating agent. The nature of the obtained product was not investigated further. It is noteworthy in this connection that stable Sn-0S02F bonds have been found in a 131 large number of Sn(IV) compounds (e.g. SnC^CSO^F^ ). The electrical conductivity of the neat liquid is higher than for selenium (IV) oxychloride. The following self-dissociation process has been invoked to explain the observed electrical 181 conductivity of the latter : —> _ -„ + 2 SeOCl2 < SeOCl + SeOCl3 By analogy, the self-dissociation process for SeOfSO^F^ might be: 2 Se0(S0_F)„ v Se0(S0„F)+ + SeO(SO-F),~ 3 Z 3 33 In view of the unsuccessful attempts to prepare KSeO(S03F)3 , it seems more appropriate to describe the self-dissociation process as: + Se0(S0TF)_ . A Se0(S0_F) + S07F~ 3 Z 3 3 Two consequences arise from the fact that both SeOCl^ and SeOCSO^F)^ are significantly dissociated in the liquid phase: (1) Both should be strong electrolytes in a protonic solvent. (2) Both may possibly undergo ligand exchange reactions via an ionic mechanism. Thus, selenium CIV) oxychlorofluorosulphate, SeOClfSO^F), is prepared by the direct addition of SeOCl2 to SeO(S03F) ';. The formation of a new compound, as opposed to a solution of the starting materials, is indicated by the observed increased viscosity of the resulting liquid. More direct evidence is provided by a Raman vibrational spectrum as is discussed in a later section. 10.3.2 Conductimetric Studies Solutions of SeOCl2 and SeO(S03F)2 in HS03F were studied by conductivity measurements in hopes of detecting the formation of oxycations in this solvent. Both solutes are miscible with HS03F and behave as bases in HS03F, as is evidenced by a continued increase in conductivity upon the addition of KS03F. The conductivity results together with the reference solute KS03F are listed in Table 33 and plotted in Figure 32. A comparison of the specific conductivities of Se0(S03F)2 with those of the reference solute, KS03F, indicate that SeO(S03F)2 behaves as a moderate base in HS03F; The simplest mode of ionization + can be formulated as in the neat liquid, SeO(S03F)2 = SeOCS03F) + S03F where Se0(S0„F) is incompletely dissociated. TABLE 33: (a) Specific Conductivities of SeOCl and SeO (S0_F) in HSO F at 25.00° SeOCl2 SeO(S03F)2 Molality K X 10 Molality K x 10 2-1-1 2 -1 -1 x 10 (ft cm ) x 10 (ft cm ) 0.000 1. 265 0.000 1.169 0.267 7. 000 0.091 3.562 0.683 21. 27 0.113 4.010 1.956 41. 24 0.159 4.167 3.747 74. 94 0.231 5.989 4.368 91. 47 0.384 8.001 5.159 106. 7 0.677 10.83 6.767 134. 5 0.986 13. 21 7.255 142. 6 1.341 15.55 7.952 154. 8 1.750 17.61 8.708 168. 2 2.308 19.82 2.517 20.62 3.162 23.01 ./continued 172 Table 33 Continued (b) Interpolated specific conductivities of SeO(SCLF) and SeOCl 2 -1 -1- 104 (ft" cm J Molality x 10 K X KS03F SeO(S03F)2 SeOCl2 0.00 1 085 1 148 1.265 1.00 25 8 13 3 27.5 1.50 38 0 16 4 36.3 2.00 50 0 18 7 45.5 2.50 62 5 20 5 54.8 3.00 72 7 22 3 64.0 3.50 84 9 24 2 73.5 4.00 97 0 82.5 IO2 Molality 174 At a given conductivity, the degree of dissociation, a, for the above equilibrium can be calculated from a = m.,,,. „/m„ n,„. , and the KSO^F SeOCSO^Fj^ equilibrium constant K^, from = "^SeOCSO F) ^ 1~a* Calculated •3 2 values of are found to vary only slightly in the concentration range _3 0.01 to 0.04 molal, and a mean value of 4.0 x 10 was obtained for K^. Higher specific conductivities obtained for the solute SeOCl^, can be accounted for by the following series of equations: + Se0Clo —' • Se0Cl + Cl" Cl" + HS03F v.\ HCl + S03F" + HS03F + HCl - -» H2C1 + S03F" Chlorides ions formed in the first step are immediately protonated by HS03F with the release of fluorosulphate ions, which account for the specific conductivities being close to those of the reference solute. HCl behaves as a weak base in HS03F, and is protonated only to a small extent. The reaction of oxidizing agents such as Se02 with selenium in ++ ++ HS03F results in the formation of the species Se^ and Seg To examine the oxidizing properties of SeOCSO^)^ in HS03F, a solution of SeO(S03F)2 in HS03F was titrated conductimetrically with elemental selenium. The results are given in Table 34 and Figure 33. The selenium additions impart a green colour to the solution ++ as the result of the formation of the species Seg . The solution of ++ Seg was found to be stable over a period of 24 hours , with only slight oxidation to the more stable species Se^++. The following reactions are involved in the titration: 175 TABLE 34: Conductimetric Titration of SeO(S03F)2 in HSC^F with Elemental Se SeO(S03F)2 Se 4 4 Molality K x 10 Molality K X 10 2-1-1 2 -1 -1 x 10 (Sl cm ) x 10 (Si cm ) 0 0000 1 148 0 0000 24 61 0 1620 4 269 0 2234 25 93 0 3671 7 514 1 575 34 97 0 7479 11 48 2 908 43 93 1 080 13 89 4 932 58 49 1 560 16 59 7 230 76 43 1 856 18 00 9 312 91 09 2 168 19 35 11.57 119.7 2 757 21 60 3 607 24 61 176 Conductivity curve of SeO(S03F)2 in HSO F at 25 °C. IO2 Molality IO2 Molality SeO(SOF) Selenium 3 2 FIGURE 33 177 ++ SeO(SO„F). + 15 Se + 2 HSCLF > 2 SeQ + 4 SO_F + H„0 j 2 o o J 2 H.O + HSO-F • H_0+ + S0-F~ • H.S0. + HF z o •< j j — Z 4 Assuming that the generation of 1 mole of ^0 in HSO^F produces 0.60 — 188 — moles of SO^F , 0.306 moles of SO^F / mole of solute are expected for the occurrence of the above processes. A comparison of the slopes of the Se and KSO^F conductivity curves are in good agreement with this expectation. While SeOCSO^F)^ is only partially dissociated in HSO^F, it is fully ionized in superacid media. The conductivity data are shown in Table 35 and Figure 34. The end point occurs when 1 mole of superacid has been neutralized by 1 mole of SeOCSO^F)^- Past the end-point, a slight conductivity decrease is observed up to 1.9 molal SeOCSO^F^- This lowering in conductivity, as opposed to an increase from the moderately basic behavior of Se0(S0_F)9 in HS0_F, is unexpected. 10.3.3 Vibrational Spectra The Raman vibrational spectra of Se0(S03F)2 and SeO(S03F)Cl are reproduced in Figures 35 and 36- The observed frequencies, relative intensities and tentative assignments are given in Table 36- Two features of the Raman spectrum of SeOfSOgF^ are immediately apparent: (1) The Se=0 stretch occurs at a frequency of 1044 cm 1, much higher than SeOF^ and SeOC^, which have their Se=0 stretching frequencies at 1012 cm 1 and 955 cm 1 respectively1^ (c.f. P=0 in 40 190 POCSO^F)^ which is lower than that observed for POF^ ). TABLE 35 SpecificiConductivities of SeOCSO^F)^ In Superacid 4 Mole Ratio K X 10 1 SeO(S03F)2/Acid fc'^m" ) 0.0000 276.5 0.0211 264.6 0.0442 251.3 0.0669 238.5 0.1065 212.6 0.1267 207.2 0.1547 193.7 0.1821 179.8 0.2083 168.9 0.2442 152.8 0.29451 134.8 0.3594 113.2 0.4167 97.72 0.4651 85.76 0.5302 72.26 0.5862 63.24 0.6353 57.01 0.7232 49.98 0.795 2 46.29 0.9127 42.61 0.9910 40.72 1.022 40.09 1.119 38.89 1.310 37.48 1.424 36.65 Conductivity Study of SeO(S03F)2 in HSbF; ( S03F)4 at 25° Base • SeO(S03F)2 o KS03F 180 181 FIGURE 36 : RAMAN SPECTRUM OF SeO(S03F)Cl 182 TABLE 36: The Raman Vibrational Spectra of Se0(S0 F) and SeO(SO F)C1 SeO(S03F)2 SeO(S03F)Cl Assignment* (cm 1) Intensity (cm-1) Intensity 1430 dep 1 1405 dep 0.2 v?(A") 1220 p 6 1212 p 1.0 1055 p 4 1071 p 1.5 .^4CA,) 1044 p 4 1048 p 0.5 Se=0 str. (SeO(S03F)2) 1011 p 2.5 Se=0 str. (SeO(S03F)Cl) 944 p 0.2 Se=0 str. (SeOCl2) 848 p 3 839 p 0.5 v2(A') 639 p 5 6635 p 1.5 v3(A«) (SeO(S03F)2) 6622 p 2.0 v3CA«) 588 p 2 587 p 0.2 v5(A') 551 dep 2 551 dep 0.1 v8(A») 452 4 Se-0 asym. str. 446 p 4 430 p 10.0 Se-0 sym. str. 411 p, sh 2.0 Se-Cl str. 390 br 1 v9(A") 340 sh 1 0=Se02 asym bending 311 p 4 0=Se02 sym. bending 276 1.0 0=Se0Cl bending 265 p 10 0=Se02 sym. bending 232 dep 6 232 dep 1.0 v6(A") 175 sh 2 171 2.0 S03F Torsion 210 1.0 0=Se0Cl bending * Symbols describing the covalent S03F group vibrational modes are taken from the correlation table on page 130 183 and (2) The asymmetric and symmetric SO^ stretching frequencies found at 1430 cm 1 and 1220 cm 1 are considerably lowered from the normal -1 72 positions of about 1500 and 1250 cm for a covalent SO^F grouping This is reminiscent of the covalent SO^F groups contained in KBrfSO^F)^ and discussed in Chapter 8, where the lowering of the SO^ stretching frequencies was attributed to polar bromine-oxygen bonds. It seems feasible that the Se-0 bonds in SeO(S03F)2 have a certain degree of ionic character, a property that is already expressed in the high conductivity of the neat liquid, i.e. • A positive charge on the selenium atom would account for the strong Se=0 bond, and the lowering of the S02 stretching frequencies from where they are found in a covalent SO^F group. The self-ionization process suggested in section 10.3.1 would produce such a charged species. The interpretation of the spectrum of Se0(S0gF)2 can be made by theory of group vibrations. As a first approximation for the 12 atom molecule where a maximum number of 30 vibrational modes are possible, assignments are made on the basis of 2 independent sets of group vibrations which are as follows: (1) the SeOY2 vibrations (where Y represents the SO^F group as a single unit), and (2) the covalent SO^F group vibrations giving rise to 2 sets of 9 vibrations; the two sets will be indistinguishable if the fluorosulphate groups are equivalent. The covalent SO^F group vibrations can be assigned as follows: The asymmetric and symmetric SO^ stretching frequencies at 1 1 1430 and 1220 cm" , the S-F stretch at 848 cm" , the S03, deformation frequencies at 639, 588 and 551 cm 1 and the S0„ rocking frequencies 184 at 390 and 232 cm account for 8 of the 9 vibrations expected for a covalent SO^F group. The remaining mode of vibration, the SeO-S stretch is most difficult to assign, and possibly coincides with the absorption peak due to the Se=0 stretch. Therefore, a highly tentative assignment for the SeO-S stretch is made at 1055 cm The Se=0 stretch is assigned to the shoulder at 1044 cm *, although one can equally make the assignments in reverse for these latter two absorptions. The SeOY^ vibrations can becassigned by comparison to bromine- 29 oxygen vibrations in KBrfSO^F^ and BrSO^F . The three stretching modes are assigned at 1044 cm * (Se=0 stretch), 446 cm * (Se^-0 symmetric stretch) and to the shoulder at 452 cm * (Se-0 asymmetric stretch). The three bending modes are assigned at 340,311 and 265 cm The absorpti band at 175 cm ^ is assigned to one of the torsion modes. The remaining torsion, modes, possibly occurring below 150 cm * were not observed due to the strong Rayleigh scattering of the laser beam in this region. The Raman vibrational spectrum of a solution from an equal number of moles of SeOCl2 and SeO(S03F)2 is consistent with the formation of SeCS03F)Cl. The replacement of a S03F group in SeO(S03F)2 by chlorine results in the following changes in band positions: Cl) The Se=0 stretching frequency is lowered to 1011 cm ^. (Shoulders, -1 at 1048 and 944 cm are due to Se0(S03F)2 and Se0Cl2 respectively, which are in equilibrium with Se(S03F)Cl). (2) A sharp intense band at 1071 cm ^ is assigned to the S-OSe stretch. This assignment is consistent with the expected increase in the S-OSe bond strength. 185 (3) The Se-0 symmetric stretch is lowered in frequency to 430 cm"^; the Se-0 asymmetric stretch can not be resolved from this intense band. A shoulder band at 411 cm * is assigned to the Se-Cl stretch. Assignments for the covalent SO^F group vibrations are made by comparison to the spectrum of SeOCSO^F),,. The remaining bands are assigned to Se-O-Cl bending and torsion modes. The Raman spectra of solutions of SeOCl2 and SeO(S03F)2 at different mole ratios are tabulated in Table 37. All bands in these spectra can be assigned to SeOCl2, SeO(S03F)2 or SeO(S03F)Cl. A concentrated solution of SeOCl2 in HS03F shows a peak at 1011 cm * , indicating the fromation of SeO(S03F)Cl, and also a weak absorption band at 2906 cm * that can be assigned to H-Cl stretching. These observations are in agreement with the conductivity results. 10.3.4 Nuclear Magnetic Resonance Studies The results of the NMR spectra are summarized in Table 38. 19 The F NMR spectra of SeO(S03F)2 and SeO(S03F)Cl show sharp single peaks at chemical shifts downfield from the corresponding sulphur 19 compounds, S0(S03F)2 and S0(S03F)C1 which have their F NMR chemical 44 shifts at -52.7 and -50.1 ppm respectively . No fine structure 19 77 or even peak broadening due to F - Se coupling was detected. 77 The Se chemical shifts for the compounds of type 0=SeXY (X,Y = Cl, S03F, F) show the expected trend for an increase in the 77 electronegativity of X or Y. The Se chemical shift for SeOF2 reported TABLE 37: Raman Spectra of the Solutions: SeOCl2 in SeO(S03F)2 and SeOCl in HS0_F (Frequencies in Cm *) Mole Ratio SeOCl2:SeO(S03F)2 Mole Ratio SeOCl2:HS03F 0:1 1:3 1:1 9:1 1:0 0:1 1:3 * 175 (2) 178 (2) 171 C2) 162 (3) 161 (4) 176 (3) 215 (3) 210 (1.5) 210 (.5) 212 (1) 232 (6) 232 (5) 232 CD 254 (2) 255 (3) 265 (10) 268 (5) 276 CD 277 (3) 279 (3) 311 (4) 311 (2) 347 (4.5) 390 (1) 390 CD 391 (10) 390 (10) 405 (8) 405(3) sh 411 (2) sh 446 C4) 436 (10) 430(10) 422 (6) 421 (7) sh 452 (4) 435 (10) 551 (2) 551 (.3) 551 CD 555 (10) 556 (1) 588 (2) 588 (.6) 587 (.2) 560 (10) 560 (1) sh 639 (5) 629 (4) 622 (2) 635 (1.5) 640 (0M) 848 (3) 852 CD 839 (.5) 850 (10) 842 (1) 911 (.2) 944 (.2) 944 (2) 944 (2) 951(.4) 968 (6) 968 (.4) 1011 (2.5) 998 CD 1011 (.4) 1044 C4) 1055 C3) 1048 (.5) 1055 C4) 1071 (1.5) 1071 (D 1088 (1) 1220 (6) 1218 C4) 1212 (1.0) 1218 (.2) 1230 (6) 1212 (1) 1430 CD 1428 CD 1405 (.2) 1445 (3) 1440 (.2) 2906 (.4) Numbers in brackets indicate relative intensity 187 by Birchall et al. is listed in Table38 for comparison. Absolution of SeOC^ and SeOCSO^F)^ in a 1:1 mole ratio shows a peak at 28 ppm consistent with the formation of SeOfSO^FjCl. A 60% solution of SeOC^ in HSO^F shows a chemical shift at 11 ppm in agreement with the equilibrium, SeOCl2 + HS03F = SeO(S03F)Cl + HCl, as suggested by the vibrational and conductivity data. TABLE 38 Chemical Shifts of some Oxy Selenium Compounds C0mP°Und 19F NMRa 77Se NMRb SeOCl2 - 0 SeOCl2 in HS03F - 11-5 Se0(S03F)Cl -47.7 28-10 SeO(S03F)2 -48.6 69-5 C SeOF2 100.6 a chemical shifts in ppm from iCFCl3 external standard b chemical shifts in ppm from SeOC^ c reference 212 10.3.5 Conclusions As mentioned briefly in the discussion of the Raman spectra of Se0(S03F)2 and Se0(S0gF)Cl, strong evidence for a strengthening of the Se=0 bonds via pff-dTr bonding exists. Since the molecular structure of selenyl compounds of the type SeOX2 and SeOXY can be approximately described as distorted tetrahedral with a stereochemically lone pair, 194 the same arguments as presented by Cruickshank for P-0 and S-0 compounds with tetrahedral coordination around P and S respectively, apply here, with the sole departure that not 3d but 4d outer orbitals are involved as ir-electron acceptor orbitals. The very same features 194 observed by Cruickshank using extensively x-ray data and by 218 Gillespie and Robinson ,using vibrational spectra as the main source of information are found for selenium-oxygen compounds. The Se-0 bond distances and the Se-0 stretching force constants show a linear relationship, both depending strongly on the electro• negativity of X and Y. -These features have been discussed and t D f *• 181,197,198,213,214 interpreted by Paetzol d .i n a serie. s or f publicationui- s ' ' . 198 It has also been noted that for compounds of the type E0X2 and EOXY (E = S,Se), a linear correlation exists between SO and SeO stretching force constants for identical X and Y ligands. This relationship can be extended to SO^, S02XY, SeO^ and Se02XY with replacement of the lone pair in the tetravalent compounds by another oxygen. The following features are noteworthy: 189 (1) Stretching force constants (and, of course, element-oxygen stretching frequencies) are higher for the type molecules 50^X^ or SeG"2X2 than for the type S0X2 or SeOX2> indicating a higher degree of p-rr-d^ bonding in the former type. (2) The stretching force constants for sulphur-oxygen compounds are much higher than the corresponding selenium-oxygen compounds. These trends are illustrated in Table 39, where a number of selected examples have been listed: TABLE 39: S-0 and Se-0 Stretching Force Constants Compound Stretching Force Constant (mdyn/A) S02F2 12.04 S0F2 11.17 S02(0CH3)2 10.4 S02 :9.91 S0C12 9.84 = S04 6.82 Se02F2 8/07 SeOF2 7.91 Se02(0CH3)2 7.36 Se02 7.53 SeOCl2 5.74 190 The second feature indicates stronger overall bonding in S-0 compounds and very likely stronger pn-idir bonding than for the corresponding 198 selenium compounds. Paetzold has found an empirical relationship between the S-0 and Se-0 stretching force constants: 2 687 f 2 30 ? fso = - - se0 " ' ; fSeO = °-593-fS0 + 1-37 Applying the above information to the case of SeOCSO^F)^ where bridging between Se(IV) and S(VI) occurs via oxygen, it seems safe to say that competition for tr-electron density exists between Se;(IV) and S(VI), and that S(VI) as the better TT acceptor in this system, will withdraw n-electron density from selenium, resulting in weaker and polar bonds. As pointed out earlier, the observed electrical conductivity indicates an ionic dissociation in the liquid state. Also, the positively polarized selenium will attract electron density from chlorine in Se(S03F)Cl resulting in the observed relatively high frequency for Vg^^* and from oxygen in the Se=0 unit, resulting in unusually high values for v CHAPTER ELEVEN CONCLUSIONS AND SUMMARY The presented experimental results are concerned with three principal groups of compounds: (a) oxyfluorides of the halogens and selenium, (b) oxyfluorosulphates of the halogens and selenium and (c) oxycation complexes of perfluoroanions with the central atoms of the cations being Cl, I and Se. All three classes of compounds can + + ++ be interpreted as containing the oxycations ClO^ , I02 and SeO or SeOX+ where X = halogen or SO^F, or as giving rise to the formation of oxycations when used as solutes in a strong protonic acid solvent such as HS03F. 215 187 With the exception of SeOF2 and SeOF^NbF^ , no structural details for any of these compounds are available at present, and therefore, the bonding in these compounds must be discussed on the structural information obtained in this study, mainly from vibrational spectroscopy and solution studies. This discussion can only be qualitative, which is the intention of the following summary of the bonding involved in the following compounds: (1) C102F - The bonding in chloryl fluoride is best described in terms 114 of the (TT* - p) a model also employed in the analogous compounds FNO and FN02- (2) ClO^O^F - Chloryl fluorosulphate is strongly dissociated as + according to ClO^O^F = C102 + SO^F , as the simplest equation to + describe the dissociation. While the presence of discrete C102 cations are likely responsible for the red colour of this compound, 192 it is believed to be covalent with a highly polar C^Cl^* - ^"OSC^F grouping. On the other hand, (ClO^^^O.^ is likely to be ionic, similar to (NC^)2^3^10^' A crystal structure determination of (CIC^)2S3O10 would be worth well attempting. (3) C10oAsF,, C10oSbF£, (ClO.KSnF, - These white solids are ZD ZD Z Z O predominantly ionic, but there is strong evidence for anion-cation interaction via fluorine bridging, possibly involving the b^ + antibonding orbitaliof C102 or outer 3d orbitals with lone pair orbitals on fluorine. The order of interaction appears to be: AsF^ <*'SbFg < SnF^ . This type of interaction has also been "found + for the corresponding complexes containing the cations, C1F2 and + + + N F + and NF + BrF2 , but not the nitrogenheterocations ko , N02 , 2 3 ^ 2 ' All three chloryl compounds appear to be good sources for solvated + C102 cafcions. (4) I02F - The structure of I02F appears to be polymeric, possibly involving both oxygen and fluorine bridging. + rou s (5) IC^AsF^ - There does not appear to be I02 g P present in + this compound as ther are C102 groups in ClC^AsF^. There is evidence + _ for a polymeric (I02 )n cation and discrete AsF^ anions. 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