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High valent element fluorosulfates and fluoride- fluorosulfates Mistry, Fred 1993

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HIGH VALENT ELEMENT FLUOROSULFATES AND FLUORIDE- FLUOROSULFATES by FRED MISTRY B. Sc. (Chemistry) University of British Columbia, 1987  A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY in THE FACULTY OF GRADUATE STUDIES CHEMISTRY  We accept this thesis as conforming to the required standard  THE UNIVERSITY OF BRITISH COLUMBIA August 1993 © Copyright Fred Mistry, 1993  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 or her representatives. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission.  Department of Chemistry The University of British Columbia 2036 Main Mall Vancouver, Canada V6T 1Z1 Date: 12 October, 1993  ii  ABSTRACT  The synthesis of the binary tetrakis(fluorosulfates) of the group 4 elements (Ti, Zr or Hf), all potential Lewis acids in fluorosulfuric acid (HSO 3 F), was carried out in order to extend the presently known conjugate superacid systems of the type HSO 3 F-M(SO 3 F) n . Metal oxidation by bis(fluorosulfuryl) peroxide (S 2 06 F2), in HSO3F, was the preferred route to these compounds. Although the polymeric M(SO 3 F) 4 compounds were insoluble in HSO 3 F, they were found to be good fluorosulfate acceptors as was demonstrated by the synthesis of their ternary salts of the type Cs 2 [M(SO 3 F) 6 ] (M = Ti, Zr or Hf). The resulting cesium salts are thermally stable up to 250°C and, as is evident from their vibrational spectra, are structurally similar to Cs 2 [M(SO 3 F) 6 ] (M = Pt, Pd, Ge and Sn) reported previously. The synthetic use of fluorine fluorosulfate (FOSO 2 F) in the preparation of mixed fluoride fluorosulfates of the type EF,(SO 3 F),„, some of them potential Lewis acids in both HF and HSO 3 F, was explored. Reactions involved the oxidative addition of FOSO 2 F to a number of binary fluorides (SF 4 , SeF4 and AsF 3 ) as well as the simultaneous fluorination and fluorosulfonation of elements (Br and I). The compounds obtained were studied by vibrational and NMR spectroscopy. Well-defined, monomeric compounds of the type EF 5 SO 3 F (E = S, Se) were formed in the reactions with SF 4 and SeF 4 . The reactions with AsF3, I 2 and Br2 yielded nonstoichiometric materials of the type EF n (SO3F)5_„ (E = As, I) and BrF n (SO 3 F) 3 _,i ; the nonstoichiometric property of these compounds has not been previously recognized. The oxidation of molybdenum pentafluoride (MoF 5 ) by either S 2 06 F 2 or FOSO 2 F produced MoOF3(SO3F) in addition to MoF 6 and S20 5 F2 rather than the expected MoF 5 SO3 F. The solvolysis of Mo2 (CH 3 C00)2 in HSO 3 F resulted in the formation of the binuclear compound Mo 2 (SO3F) 4 , the only binary fluorosulfate of molybdenum known so far. Addition of acetonitrile (CH 3 CN) to this compound yielded Mo(SO 3 F) 2 (CH 3 CN) 4 . The fluorosulfate acceptor ability of MoF 5 was demonstrated by the formation of Cs[MoF 5 SO 3 F] in a solid-state reaction between MoF 5 and CsSO3F.  111  The nucleophilicities of various fluoro- and oxy-acid anions in (CH 3 ) 2 SnX 2 salts have been studied in the past using M8ssbauer spectroscopy. All the examples studied had hexacoordinated tin, bidentate symmetrically bridging X-groups, and a linear C-Sn-C grouping. An attempt was made to extend the linear plot of the isomer shift (6) vs. the quadrupole coupling (SEQ),both of which correlate well with anion basicities, to dimethyltin(IV) carboxylates. To this end, the crystal structures of (CH 3 )2Sn(OOCR)2 (R = H, CF 3 and CH 3 ) were solved. The formate salt, (CH 3 ) 2Sn(OOCH) 2 , is a sheet polymer with symmetrically bridging formate groups and a linear C-Sn-C group, and fits the correlation well. The trifluoroacetate group is unsymmetrically bridging in (CH 3 ) 2 Sn(00CF 3 ) 2 and the acetate group is chelating in (CH 3 ) 2 Sn(OOCH 3 ) 2 . In both these compounds, the C-Sn-C moiety shows significant departure from linearity; this is reflected in their MOssbauer data which do not fall on the linear plot of S vs. SEQ . It appears that for 00CR where R^H, the steric hindrance is increased to the point where a regular bridging structure cannot be obtained, and results in a distortion about the tin atom.  iv  TABLE OF CONTENTS  ABSTRACT ^  ii  TABLE OF CONTENTS ^  iv  LIST OF TABLES ^  viii  LIST OF FIGURES ^  x  LIST OF SYMBOLS ^  xii  ACKNOWLEDGEMENTS ^  xiii  CHAPTER 1. INTRODUCTION ACID-BASE CONCEPTS ^  1  SUPERACIDS ^  4  FLUOROSULFURIC ACID ^  8  BIS(FLUOROSULFURYL) PEROXIDE ^  11  THE FLUOROSULFATES ^  14  Synthetic Routes to the Fluorosulfates ^  15  Decomposition of the Fluorosulfates ^  23  STRUCTURAL ANALYSIS OF THE FLUOROSULFATES ^ 23 Structures of some fluorosulfates ^  24  Vibrational analysis of the fluorosulfate group ^  28  SUMMARY ^  31  CHAPTER 2. EXPERIMENTAL GENERAL EQUIPMENT ^  38  Utility Vacuum Lines and Glassware ^  38  Ampoule Key ^  41  Sublimation Apparatus ^  41  Vacuum Filtration Apparatus ^  41  v  Metal Vacuum Line and Vessels ^  41  Inert Atmosphere Box ^  44  INSTRUMENTATION ^  44  Infrared Spectrometers ^  44  Raman Spectrometers ^  47  Nuclear Magnetic Resonance Spectrometers ^  47  Electron Spin Resonance Spectrometer ^  47  MOssbauer Spectrometer ^  48  X-Ray Diffractometer ^  49  Mass Spectrometer ^  49  Melting point Determinations ^  49  MICROANALYSES ^  49  GENERAL REAGENTS ^  50  SYNTHESIS OF BIS(FLUOROSULFURYL) PEROXIDE (S 2O 6 F2) AND FLUORINE FLUOROSULFATE (FOSO 2F) ^ 50 The Flow System ^  51  The Flow Reactor and Catalyst ^  51  Product Isolation and Purification ^  54  CHAPTER 3. DIMETHYLTIN(IV) CARBOXYLATES AND ANION BASICITIES INTRODUCTION ^  57  Organotin compounds ^  58  MOssbauer Spectroscopy ^  62  EXPERIMENTAL ^ Reagents ^  72 72  DISCUSSION ^  73  CONCLUSIONS ^  85  vi  CHAPTER 4. FLUOROSULFATE DERIVATIVES OF GROUP 4 ELEMENTS INTRODUCTION ^  88  EXPERIMENTAL ^  90  Reagents ^  90  Synthetic reactions ^  91  DISCUSSION ^ CONCLUSIONS ^  94 106  CHAPTER 5. FLUOROSULFATE DERIVATIVES OF MOLYBDENUM INTRODUCTION ^  111  EXPERIMENTAL ^  114  Reagents ^  114  Syntheses ^  115  DISCUSSION ^  116  CONCLUSIONS ^  130  CHAPTER 6. SELECTED REACTIONS OF FLUORINE FLUOROSULFATE INTRODUCTION ^  134  EXPERIMENTAL ^  140  Reagents ^  140  Reaction Vessels ^  141  Synthetic Reactions ^  142  DISCUSSION ^ The stoichiometric fluoride- fluorosulfates ^  143 145  The non-stoichiometric fluoride- fluorosulfates ^ 152 The decomposition of fluorine fluorosulfate ^ CONCLUSIONS ^  158 160  vii  CHAPTER 7. CONCLUSIONS GENERAL COMMENTS ^  165  SUMMARY AND SUGGESTIONS FOR FUTURE WORK ^ 166 CONCLUDING REMARKS ^  168  APPENDIX A. THE GENERATION OF Au(II) IONS INTRODUCTION ^  169  General Comments ^  169  ESR Spectroscopy in the study of Au 2+ ions ^  174  EXPERIMENTAL ^  176  Reagents ^  176  Synthetic Reactions ^  177  DISCUSSION ^  178  CONCLUSIONS ^  184  viii  LIST OF TABLES  Table 1.1. Some Physical Properties of HSO 3 F, HF, H 2 O and H2SO4 ^ 10 Table 1.2. Transition Metal Fluorosulfates obtained via solvolysis in HSO 3 F ^ 16 Table 1.3. Transition Metal Fluorosulfates obtained via oxidation with S 2 06 F2 in the presence or absence of HSO 3 F ^ 18 Table 1.4. Miscellaneous Routes to the Transition Metal Fluorosulfates ^ 22 Table 3.1. 119 Sn MOssbauer data of some dimethyltin(IV) derivatives of strong protonic acids and superacids ^  70  Table 3.2. Crystallographic data for (CH3)2Sn(O2CR)2 derivatives ^ 75 Table 3.3. Bond lengths (A) and angles (°) for (CH 3 ) 2 Sn(O 2 CH)2 with estimated standard deviations in parentheses ^ 77 Table 3.4. Bond lengths (A) and angles (°) for (CH 3 ) 2 Sn(O 2 CCH 3 ) 2 with estimated standard deviations in parentheses ^ 79 Table 3.5. Bond lengths (A) and angles (°) for RCH 3 ) 2 Sn(0 2 CCF 3 ) 2.1 11 with estimated standard deviations in parentheses ^ 82 Table 4.1. Binary and Ternary Metal Fluorosulfate Derivatives obtained by Oxidation with S 206 F 2 in HSO 3 F ^  97  Table 4.2. Infrared Spectra of [Au(SO 3 F) 3 ] 2 and the Tetrakis(fluorosulfates) of Zirconium, Hafnium, Tin, Platinum and Related Compounds ^ 101 Table 4.3. Vibrational Spectra of Cs 2 [M(SO3 F) 6 ], with M = Ti, Zr, Hf and Sn ^ 104 Table 5.1. Vibrational data for MoOF 3 (SO 3 F) ^  118  Table 5.2. Vibrational frequencies of MoOF 4 , MoOF 3 (SO 3 F), and MoO(SO 3 F) 4 between 1500 and 550 cm -1 ^ Table 5.3. 19 F NMR data for MoOF 3 (SO 3 F) ^  121 121  Table 5.4. Vibrational data for Cs[MoF 5 S0 3 F] along with selected Raman bands of [AsF 5 (SO 3 F)] -  ^  124  ix Table 5.5. Infrared Spectrum of MoF s •CH 3 CN ^  126  Table 5.6. Infrared data for Mo2(SO3F)4 ^  127  Table 5.7. Vibrational Spectra of Mo(SO3F)2(CH3CN) 4 ^  129  Table 6.1. Some physical properties of FOSO 2F and S 2 06 F 2 ^  135  Table 6.2. Summary of the attempted reactions of fluorine fluorosulfate, FOSO 2 F ^  144  Table 6.3. Vibrational Spectra of the Pentafluorofluorosulfates of Sulfur, Selenium and Tellurium in the Frequency Range of 1500-110 cm -1 ^  147  Table 6.4. Raman Spectra of Br(SO 3 F) 3 , BrF n (SO 3 F) 3 _ n and IF 6 (SO 3 F) 6 _„ ^  153  Table 6.5. Characterization of the fluoride-fluorosulfates of sulfur, selenium, arsenic, bromine and iodine by NMR ^ Table A.1. Ionization potentials for group 11 elements ^  156 169  x  LIST OF FIGURES  Fig. 1-1. Crystal structure of gold tris(fluososulfate), [Au(SO3F)3J 2 ^ 26 Fig. 1-2. Vibrational frequencies of the fluorosulfate group ^  30  Fig. 2-1. Vacuum line and trap used for manipulation of volatiles and liquids ^ 39 Fig. 2-2. Reaction vessels used in synthesis ^  40  Fig. 2-3. a) Ampoule key used for opening and resealing glass vials under vacuum b) Sublimation apparatus used for purification of air-sensitive compounds ^ 42 Fig. 2-4. Filter apparatus for filtration of air-sensitive samples ^ 43 Fig. 2-5. Kel-F reaction vessel for reactions in HF ^  45  Fig. 2-6. Gas cell for obtaining infrared spectra on gaseous samples ^ 46 Fig. 2-7. Flow system used in the preparation of S 2 0 6 F 2 ^  53  Fig. 2-8. Flow reactor used for the synthesis of bis(fluorosulfuryl) peroxide, S20 6 F2 -54 Fig. 3-1. Crystal structure of (CH 3 ) 2Sn(SO 3 F)2  61  Fig. 3-2. a) Decay scheme for 119 "15n b) Block Diagram of MOssbauer Spectrometer ^  63  Fig. 3-3. Diagrammatic scheme showing the isomer shift ^  65  Fig. 3-4. Diagrammatic scheme for quadrupole splitting ^  67  Fig. 3-5. Correlation of isomer shift and quadrupole splitting for dimethyltin(IV) salts of strong fluoro- and oxyacids ^ 71 Fig. 3-6. a) Stereoview down the  a axis of the structure of [(CH3 ) 2 Sn(02 CH) 2],  b) Packing diagram of [(CH 3 ) 2 Sn(0 2CH) 21 n viewed along the a axis ^ 76 Fig. 3-7. a) Stereoview of the structure of (CH 3 ) 2 Sn(O 2 CCH 3 ) 2 b) Packing diagram of (CH 3 ) 2 Sn(O 2 CCH 3 ) 2 viewed along the b axis ^ 78 Fig. 3-8. a) Stereoview of the structure of [(CH 3 ) 2 Sn(0 2 CCF 3 ) 2],, b) Packing diagram of RCH 3 ) 2 Sn(0 2 CCF 3 ) 2L viewed down the  a axis ^ 81  Fig. 4-1. Infrared Spectra of solid Hf(SO 3 F) 4 and Zr(SO3F) 4 ^  100  xi Fig. 4-2. Infrared Spectra of Cs2[Ti(SO3F)6] and Cs 2 [Hf(SO 3 F) 6 ]  105  Fig. 5-1. Raman spectra of Mo0F 3 (SO 3 F)  119  Fig. 5-2. a) 19 F NMR of MoF6 b) 95 Mo NMR of MoF 6  123  ^  Fig. 5-3. 19 F NMR of polysulfuryl fluorides formed in the reaction between SO 3 and MoF 6 125 Fig. 6-1. Gas phase infrared Spectra of SF 5 SO 3 F and SeF5SO3F ^ 149 Fig. 6-2. Raman Spectrum of liquid SF 5 SO3F ^  150  Fig. 6-3. Decomposition scheme of FOSO 2F ^  159  Fig. A-1. Structure of Au[AuC1 4] ^  170  Fig. A-2. The coordination environment of Au(III) in [Au(SO 3 F) 3 ] 2  ^ 172  Fig. A-3. X-band ESR spectrum of Au(SO 3 F) 3 _x recorded at 100 K. ^ 179 Fig. A-4. X-band ESR spectrum of Au 2+ (solv) in HSO 3 F recorded at 100 K ^ 182  xii LIST OF SYMBOLS  b^broad (infrared, Raman) d^days (reaction time) 6^deformational vibration (infrared, Raman), chemical shift in ppm (NMR), isomer shift in mm-s -1 (MOssbauer) AEQ^quadrupole splitting in mm.s -1 (MOssbauer) ESR^electron spin resonance h^hours (reaction time) Int^intensity (vibrational spectra) J^coupling constant in Hz (NMR) K^equilibrium constant m^medium intensity (infrared, Raman), minutes (reaction time) NMR^nuclear magnetic resonance P  rocking vibration (infrared, Raman), polarization ratios (Raman)  s^strong intensity (infrared, Raman) sh^shoulder (infrared, Raman) T^temperature 7^torsional vibration (infrared, Raman) ^  vibrational frequency in cm -1 (infrared)  Ay^vibrational frequency in cm -1 (Raman)  ^  stretching vibration (infrared, Raman)  vas^asymmetric stretching vibration (infrared, Raman) vsym^symmetric stretching vibration (infrared, Raman) w^weak intensity (infrared, Raman)  xl"  ACKNOWLEDGEMENTS  I would like to thank Professor F. Aubke for his guidance and encouragement throughout the duration of this project, and Professor Helge Willner for passing on his expertise and unique brand of experimental technique to me. I would like to thank the late Dr. J.G. Hooley for introducing me to academic research, and for encouraging me to continue in the field. Professor J. Trotter is thanked for all his suggestions and help. The work on gold compounds was carried out in collaboration with G. Hwang, and I would like to thank her for letting me include some unpublished details in this thesis. Funding support from NSERC and NATO are gratefully acknowledged. The help of all the people in microanalytical, X-ray crystallography, and mass spectrometry services and the mechanical, electronic and glass-blowing workshops is much appreciated and acknowledged. I have also been assisted by numerous people in the Chemistry Department on many occasions over the years. I have tried to thank them all along the way; to those I didn't, let me thank them now.  xiv  ... a scientist must also be absolutely like a child. If he sees a thing, he must say that he sees it, whether it was what he thought he was going to see or not. See first, think later, then test. But always see first. Otherwise you will only see what you were expecting. Most scientists forget that.... You can't possibly be a scientist if you mind people thinking you're a fool.... Let people think I'm a fool. That allows me to say what I see when I see it.  Wonko the Sane  in So Long And Thanks for All the Fish by Douglas Adams  ^  1 CHAPTER 1. INTRODUCTION  General definitions and basic concepts used in this thesis will be discussed in the first chapter together with a formulation of the research objectives. The general overview includes a review of the literature in this area and an introduction to some of the methods and concepts that are of relevance to this study.  ACID-BASE CONCEPTS Descriptions of acids and bases go back to antiquity. In chemical science, there have been several interrelated classical concepts and definitions which represent a historical development where the acid-base concept is expanded to allow for more and more general usage. The first generally useful concept dates back to 1887, when Arrhenius [1] defined acids as substances containing a proton H + , and bases as substances containing the hydroxyl ion OW. In pure water both ions can be liberated as shown by two simple examples:  HC1 (aq) ^ > 1-1 + (aq) +^(aq) ^ NaOH (aq) ^ > Na + (aq) + OW (aq) -2HC1 + NaOH ^ > NaCI + H 2 O (aq)^-3-  The third equation represents a neutralization or salt-formation reaction. The Arrhenius concept is closely connected with the then emerging view of ionic dissociation in water, ie.  H 2 O ^ H+ (aq) +^(aq)  ^-4-  The obvious limitation of this concept is the restriction to water as the solvent. The BrOnsted concept [2], developed in 1923, provides an extension to non-aqueous protonic solvents and focusses attention on the proton transfer between solvent and solute.  ^  2  This ion transfer generates acid-base pairs that differ only by a proton as is illustrated by three examples of ionic equilibria: Kid HF HF^H2F+ F ^acidl  -5-  -  basel^acid2 base2  This equation describes proton transfer between solvent molecules and is termed "ionic selfdissociation". It is quantitatively measured in terms of the equilibrium constant  Kid  which is  temperature dependent. With HF as a solvent, the solute H 2 O behaves as a base, or proton acceptor, according to:  HF + H 2 O ^ acid 1 basel  ^ H30+ + F -  -6-  acid2 base2  Equilibrium 6 also adequately describes the reverse situation with H 2 O as the solvent and HF as the solute, and it becomes apparent that of the two protonic solvents, HF and H 2 O, the former is a stronger proton donor, and the latter is a better proton acceptor. Generally speaking, knowledge of the proton donor/acceptor ability of a protonic solvent is required to designate the formulation of BrOnsted acid-base pairs. In HF as the solvent, H 2 O behaves as a base; in liquid ammonia, it is an acid, according to:  H2O + NH 3  ---'- NH 4 + + 0H -^-7-  acidl^basel^acid2^base2  Likewise, acetic acid is an acid in H 2 O but a base in H 2 SO4 . Equilibrium 7 makes the point that the classification of compounds as acids or bases is solvent dependent. An important  3  consequence of this solvent dependency is the "levelling effect" of a solvent. In a protonic solvent HA, the self-dissociation ions, H 2 A + and A - , are the strongest acids and bases in this solvent; stronger proton donors or better proton acceptors are "levelled" to these two ions. For example, dissolution of H 2SO4, HC1O4, HC1 or CH 3 COOH in H 2O will result in the formation of H 3 0 + , albeit at different concentrations, reflecting their relative proton donor ability. The solvent system concept introduced by Cady and Elsey [3] and subsequently popularized by Gutmann [4] allows the acid-base concept to be further extended to ionizing non-protonic solvents. In analogy to the BrOnstecl system which has been based on proton transfer to give ions, the solvent system concept also includes ion transfer by ions in addition to the proton. Three simple equilibria may be invoked to illustrate the solvent system concept. 2 HF^. H 2 F + + F -^ -82 BrF 3 ,^ '^ BrF 2 + + BrF4^ -92 SO2 ^' S02+ + S03 2  ^ -  -10-  The ions transferred are H + , F - and 0 2- , and the acidium ions are H 2 F + , BrF2 + and S0 2+ , while the base ions are F - , BrF 4 - and S0 3 2- . In this system, an acid is defined as any solute that will cause an increase in the acidium ion concentration, and a base as any substance that will cause an increase in the base ion concentration. A solute that will not change both ion concentrations is called a non-electrolyte. An obvious consequence of the equilibria above is that a decrease in base ion concentration implies an increase in acidium ion concentration and  vice versa. In the BrF 3 solvent system, both ions formed (BrF 2 + and BrF4) are realistic ions that occur in known compounds (eg. BrF 2 SbF 6 and KBrF 4 ). Their addition to liquid BrF 3 would increase the acidium or base ion concentrations respectively. They are termed as solvo-acids or solvo-bases. On the other hand, addition of SbF 5 or KF would have the same effect even  4  though neither one contains the acidium or base ion, but gives rise to their formation by reactions with the solvent:  SbF5 + BrF 3 ,^ '^ BrF2 + (s01 , ) + SbF6 - (01 ,) KF + BrF 3 ,-=-----  K+(sole)  + BrF4 - (soi v)^-12-  Hence they are termed as ansolvo-acids or ansolvo-bases. In order for the solvent system concept to be applicable, there must be physical and chemical evidence for the postulated ions. Sizeable and reproducible electrical conductivities provide physical evidence and the existence of compounds as discussed above for BrF 3 provides chemical evidence. On both accounts, SO 2 may not be a good instance to apply the solvent system concept: a) conductivities are small and possibly due to trace impurities (eg. moisture), and b) while S0 3 2- does exist, as for example in Na2SO3, any evidence for SO 2-1- is lacking. To rationalize the solution behaviour of SO 2 , an additional concept must be introduced. Lewis [5] defined acids and bases in terms of valence electron redistribution. In the Lewis concept, an acid is defined as an electron pair acceptor (Fe 3+ or BF3), and a base as an unshared electron pair donor (SO 2 , NH 3 or F -), thus extending the acid-base concept to nonprotonic solvents. Lewis acids and bases may be molecular (BF 3 or NH 3 ) or monoatomic (Fe 3+ or F - ). An acid-base interaction may lead to adduct-formation:  NH3 + BF 3  > NH3 BF3 -  -13-  SUPERACIDS In 1927, Hall and Conant [6] coined the term "superacid" to describe strong protonic acids that could protonate weak bases such as organic carbonyl compounds. Just as acids may be further divided as mentioned above, so can superacids which may be divided into two types, each of  5  which is explained by either the BrOnsted or the Lewis concept. Hence, we have BrOnsted superacids defined by Gillespie [7][8] as those acids that are stronger than 100% H2SO  4  (eg. HC10 4 , HSO 3 CF 3 , HSO 3 C1, HF, HSO3F), and Lewis superacids defined by Olah et al. [9)411] as those stronger than AlC1 3 (eg. SbF 5 , AsF5 , TaF5 , NbF5 ). Both these definitions, though arbitrary, are widely accepted. There are two ways of enhancing the acidity of a BrOnsted acid. One may increase the concentration of the acidium ion either by direct protonation or by decreasing the concentration of the base ion, thus shifting the self-dissociation equilibrium to the right. The concentration of the acidium ion may be increased by the addition of a strong BrOnsted acid thus inducing protonation, according to:  HF + HSO 3 F ^ , '. H 2 F + + SO3 F -^-14-  However, this method has limitations in the regime of strong acids such as those in equation (14) in cases where the two acids may have comparable acid strengths. The concentration of the base ion may be decreased by adding a Lewis acid to the Br6nsted acid. This combination gives rise to a conjugate BrOnsted-Lewis Superacid system. Thus in order to increase the concentration of H 2 F + , we may add BF3 and obtain the following:  2 HF + BF 3  ^ '' H 2F +  + BF 4 -^-15-  Conjugate BrOnsted-Lewis superacid systems are the systems of importance in the regime of high acidity. The most widely studied systems are the HF-SbF 5 and HSO 3 F-SbF 5 systems [9]-[11][12]. In the HSO3F-SbF5 system, fluoride-fluorosulfate derivatives of Sb(V) are formed in situ due to the presence of SO 3 [13]. The pentavalent antimony fluoridefluorosulfate species are stronger acids in HSO3F than is antimony pentafluoride, SbF 5 . The strongest acid in HSO 3 F so far appears to be SbF 2 (SO 3 F) 3 , although it has only been assumed  6  to exist in situ [13]. In these systems involving fluoride-fluorosulfates in HSO 3 F, an increase in the SO 3 F vs. F content results in a corresponding increase in acidity; one would expect the highest acidity for Sb(SO3F)5. However, the only species isolated so far have been identified as SbF 4 (SO 3 F) [14][15], SbF3(SO 3 F) 2 [15] and Sb2F9SO3F [15]. In the HSO 3 F-SbF5 system, dissociation of the Sb(V) fluorosulfate derivatives in situ leads to the formation of SO 3 . The presence of SO 3 in the system often leads to oxidative side reactions towards moieties such as polyhalogen cations [16][17] or certain carbocations [18]. The solvent system concept is of great use in explaining the conjugate BrOnstecl-Lewis acid. The Lewis acid is viewed as an ansolvo acid, and its effect on the BrOnsted acid systems is the removal of the base ion by complexation, resulting in an increase of the acidium ion. Solid Superacids are systems having strong BrOnsted or Lewis acidity, and may also involve  conjugate systems. The definition is of practical importance; while these media are excellent in facile product isolation, they are not encountered in this study and will not be discussed in detail. Solid superacids may be obtained by chemically or physically attaching a liquid superacid to a solid substrate having little or no acidity. This can be achieved by deposition or intercalation of materials containing acidic sites (BrOnsted or Lewis) onto the support. Another method used to increase the intrinsic acidity of a solid acid is the complexation of a suitable additional Lewis acid. The achievable acidity of conjugate BrOnsted-Lewis systems depends on the basicity of the anion, X' or EX 0+1 ; in the following equations:  2 HX EX n + X -  H2X + + X EX(n + 1)  -  -  -16-17-  In order to measure the anion basicity of strong fluoro- or oxo- protonic acids, the linear (CH3) 2 Sn 2+ cation has been used as a MOssbauer probe, and an empirical relationship  7  between the MOssbauer parameters (5 and AE,Q ) and anion basicity was reported [19]. This relationship is extended to the regime of the weaker protonic acids such as the carboxylic acids in Chapter 3, and its validity is tested by using single crystal X-ray diffraction data to explain the deviations observed for some of the compounds examined. Superacid systems based on fluorosulfuric acid as a BrOnsted acid, and high valent element fluorosulfates as Lewis acids, have been studied in the past by our group. In these systems, the binary fluorosulfates have been derived from the noble metals (Au and Pt) and the group 5 metals (Nb and Ta). The inherent problems associated with these systems are the oxidizing ability of Au(SO 3 F) 3 and Pt(SO 3 F) 4 , and the instability of Nb(SO 3 F) 5 and Ta(SO 3 F) 5 towards dissociation to give fluoride fluorosulfate derivatives (vide infra). This research is motivated, in part, by the quest for potential Lewis acids of the type E(SO 3 F) n . To fulfil this role, the compounds must be reasonably thermally stable, and sufficiently soluble in HSO 3 F; in addition, they must not act as oxidizing agents. The fluorosulfates of the group 4 elements are potentially suitable candidates since the elements easily attain the +4 oxidation state and are known to expand their coordination sphere to octahedral (and possibly higher) coordination. Their halides are known to form halo-anions complexes, and this behaviour may extend to their fluorosulfates as well. The synthesis and characterization of the high valent binary fluorosulfates of group 4 are discussed in Chapter 4. The fluorosulfate acceptor ability of these compounds is of importance and will be demonstrated by the synthesis and isolation of the ternary cesium salts of the type Cs2[M(SO3F)6] (M = Ti, Zr and HO. The pursuit of Lewis acids based on molybdenum remains a synthetic challenge, and the potential of MoF 5 and molybdenum fluoride-fluorosulfates as Lewis acids is explored in Chapter 5. In cases where binary fluorosulfates are not obtainable, or are thermally labile, emphasis can be placed on the corresponding fluoride-fluorosulfates, as discussed for antimony  (vide supra). A synthetic approach to the fluoride-fluorosulfates of some main group elements  8  is explored in some detail, along with the merits of FOSO 2F as a synthetic reagent, in Chapter 6.  FLUOROSULFURIC ACID Fluorosulfuric acid (HSO 3 F) has been known since 1892 [20] when Thorpe and Kirman first synthesized it by the reaction between hydrogen fluoride and sulfur trioxide according to:  HF + SO 3 ^ > HSO 3 F^  -18-  and is considered to be one of the strongest simple Briinsted acids. Fluorosulfuric acid is a clear, colorless liquid that is prepared commercially by the above method. Commercially available technical grade HSO 3 F can be further purified by double distillation at atmospheric pressure in a counterstream of dried nitrogen [21], and the purified product may be stored in moisture free glass containers at room temperature. HSO3F can be handled in borosilicate glasses like Pyrex without causing etching by the liquid; this facilitates its use in synthesis compared to HF which can only be handled in metal or plastic equipment, both of which are usually opaque. It should be noted, however, that any HF present in technical grade HSO 3 F will etch glass above the liquid level. Although ionic salts such as KSO 3 F have been recrystallized from their neutral aqueous solutions [22], fluorosulfuric acid itself is hydrolytically unstable in water and undergoes spontaneous hydrolysis according to:  HSO 3 F + H 2 O ^ > HF + H 2 SO 4^-19-  and hence is incompatible with aqueous systems [23]-(26], although there have been some reports on aqueous solutions of HSO 3 F [27][28]. The water that is sometimes produced in very small amounts as a reaction product acts as a base and is protonated by HSO3F:  ^ ^  9  HSO3 F + H 2 O ^ > H 3 0 + + SO 3 F -^-20-  The hydrolytic instability of the S-F bond is also manifested in most fluorosulfates, which react rapidly with moisture. At this point, the importance of using moisture free equipment and dry reagents must be stressed in order to minimize the formation of HF which causes etching of glassware in the form of a continuous hydrolytic cycle involving Si0 2 , the main constituent of glass, according to equations (19) and (21):  4HF + Si0 2 ^ > SiF 4 + 2H 2 0^ -21-  The extensive study of HSO 3 F as an ionizing, non-aqueous solvent is reflected in a number of reviews which cover many aspects of its properties and uses [29]-[32]. Its use as a fluorinating reagent when reacted with various oxides and salts of oxyacids has been discussed by Engelbrecht [33]. The broad liquid range (-89 to 163°C) allows reactions over a wide temperature range. The sufficiently high vapour pressure of HSO 3 F at room temperature makes it convenient to handle and distill the acid in vacuo. The presence of one mole of hydrogen per mole of HSO 3 F as compared to H 2 SO4 which has two moles of hydrogen per mole H2 SO4 results in less extensive hydrogen bonding in fluorosulfuric acid as compared to sulfuric acid; this is reflected in the lower boiling point and viscosity of HSO 3 F. The relatively low viscosity of HSO 3 F as compared to that of H 2 SO4 results in an increased mobility of electrolytes in the acid and also permits product isolation by filtration. HSO 3 F (mp = -89°C) is a convenient solvent for low-temperature, multi-nuclear ( 1 H, 19 F, 13 C, etc.) NMR studies. Some of the properties of HSO3 F are listed in Table 1.1 along with those of H 2 O, HF and H2SO4.  10  Table 1.1: Some Physical Properties of HSO 3 F, HF, H 2 O and H 2 SO4  HSO3F  HF  H2O  H2S O4  Molecular Weight (g•mo1 -1 )  100.07  20.00  18.02  98.08  Melting Point (°C)  -89.0  -89.4  0.0  10.4  Boiling Point (°C)  162.7  19.5  100.0  290.3  Liquid Range (°C)  252  109  100  280  Density * (g•mL -1 )  1.726  1.002  1.00 (4)  1.8269  Viscosity * (cPoise)  1.56  0.256 (0)  0.8904  24.54  Dielectric constant *  ca. 120  84  78.5  100  Autoprotolysis constant *  3.8x10-8  ca. 2x10 -12  10-14  2.7x104  *at 25°C unless otherwise mentioned.  Two important differences between sulfuric acid and fluorosulfuric acid are observed in their dissociative reactions and hydrogen-bonding. Sulfuric acid undergoes self-ionization as well as dissociation; a combination of both gives rise to a series of equilibria as follows [34]:  Self-ionization or ionic dissociation: 2 H 2 SO4 ,^  ^H3S 04 ± + H S 04  -  -22-  Dissociation: H2SO4 ..,----^ H 2 O + SO 3 `  -23-  followed by  H 2 SO4 + H 2 O ,^ H30+ + HSO 4 -  -24-  and  H 2 SO 4 + SO 3^H2S207  -25-  followed by Overall  H 2 SO 4 + H 2 S20 7^H3SO4+ + HS 2 0 7 2 H 2 SO 4 ..^ H30+ + HS 2 0 7 -  -26-27-  ^  11  Hence 100% H2SO 4 will contain the following species: neutral: H 2 SO4 (bulk), H 2S 2 0 7 and H 2 O (trace quantities) cationic: H 3 SO4 + and H 3 0 + anionic: HSO4 and HS 2 07 - . Fluorosulfuric acid undergoes the two autodissociative reactions listed below:  , H2S03F+ + SO 3 F -^K=3.8x10-9mo12kg2, 25°C^-282 HSO 3 F ^  HSO3 F ,^ '^ HF + SO3^K<3x10-7, 25°C^-29-  The latter equilibrium is the reverse of the formation reaction, and although it plays a role at temperatures above 25°C, it does not interfere with conductivity measurements at room temperature because the concentrations of HF and SO 3 are negligible. The self-ionization equilibrium in equation 28 is not believed to interfere with the interpretation of cryoscopic data in systems involving fluorosulfuric acid [29]. Fluorosulfuric acid ranks among the strongest simple protonic acids; its acidity is higher than that of H 2 SO 4 , H 2 S 2 0 7 , HSO 3 C1 and HSO 3 CF 3 , and hence is widely used in superacid systems. The Ho values of HF and HSO 3 F are of the same order (-15.1), but in practice HF exhibits lower acidity (-H0 ) because it cannot be completely freed of H 2 O which acts as a base in anhydrous HF, giving rise to F .- ions. HSO 3 F is often used in conjunction with a Lewis acid as part of a conjugate BrOnsted-Lewis Superacid system.  BIS(FLUOROSULFURYL) PEROXIDE Bis(fluorosulfuryl) peroxide (S 2 06 F2) was first prepared in 1956 as a byproduct in the synthesis of fluorine fluorosulfate by Dudley et al. [35]. In 1957, Dudley and Cady [36] reported the synthesis of S 2 06 F 2 according to the reaction:  12  AgF 2 , 160°C 2 SO3 + F2 ^ > S206F2  -30-  The first appearance of S 2 06 F 2 in the literature was in 1955, when Wannagat and Mennicken [37] reported obtaining a product of the approximate composition S 2 06 F 2 by passing an electric discharge through a mixture of SOF 2 and 0 2 in a static system at -50 to -60°C. The physical properties of this product, however, differ from those of the S 2 0 6 F 2 obtained by Dudley and Cady. Other methods for the synthesis of S 2 06 F2 include the static fluorination of SO 3 at 170°C in a nickel reactor [38], electrolysis of a solution of an alkali metal fluorosulfate in fluorosulfuric acid [39], and the photolysis of a mixture of SO 3 and F2 at 365 nm [40]. The use of NiF 2 and CuF 2 in the catalytic fluorination of SO 3 has also been reported [39]. The more exotic routes to S 2 0 6 F2 involve the xenon fluorides. The solvolysis of the xenon fluorides in fluorosulfuric acid may be used to prepare S 2 O6 F2 in very high purity [41]-[44] according to:  -75°C^25°C XeF 2 + 2 HSO 3 F ^ > Xe(SO3F)2 ^ - 2 HF  > Xe S206F2  -31-  XeF 4 + 4 HSO 3 F ^ > Xe(SO 3 F) 2 + 4 HF + S 2 O6 F 2^-3280°C > 6 XeF 5 SO 3 F ^ > Xe + 5 XeF 6 + 3 S 2 O6 F 2 -33-  6 XeF 6 + 6 HSO 3 F -6 HF  The insertion of SO 3 into the Xe-F bond of XeF 6 has also been reported to give S 2 O6 F2 [44].  XeF 6 + 6 SO 3  70°C  ^>  Xe + 3 S 2 0 6 F 2^-34-  The catalytic (AgF 2 ) fluorination of SO 3 has been adapted by our group to yield large  (ca. 1-2 kg) quantities in relatively short periods of time, as is outlined in Chapter 2. The  13  mechanism of the reaction has been discussed in detail by Wechsberg and Cady [45] as well as Leung and Aubke [46]. Bis(fluorosulfuryl) peroxide is a clear pale-yellow liquid at room temperature and, unlike the unstable peroxide S 20 6 (CF 3 ) 2 [47], may be conveniently stored at room temperature in sealed glass ampoules after careful purification (see Chapter 2). The liquid range of S 2 06 F 2 (-55 to 67°C) allows its use in chemical reactions near room temperature. Bis(fluorosulfuryl) peroxide is miscible in HSO 3 F in all concentrations, wherein it acts as a weak base [48]. Its working range may be extended up to 140°C by its solution in fluorosulfuric acid. Its chemistry has been comprehensively reviewed, more recently by De Marco and Shreeve [49]. The molecule is a dimer of SO 3 F . bridged by a weak peroxide linkage. Its well-documented characterization via vapour pressure curves [50], 19 F NMR spectroscopy [51] and vibrational spectroscopy [36][52] has been instrumental in identifying its presence using these techniques, in particular the 800 cm -1 band (p0-0) in the Raman spectrum and the singlet at 40 ppm in the  19 F  NMR. Based on vibrational spectra, the non-  planar structure of bis(fluorosulfuryl) peroxide is considered to belong to the point group C2 (similar to H202 ). The weak 0-0 link (ca. 95 kJmo1 -1 ) is responsible for the reversible equilibrium:  S206F2 ^  ^2 SO 3 F .^-35-  The SO 3 F . radicals are dark brown in the vapor phase, and are clearly observed when the temperature is raised to ca. 60 ° C or higher. The pale yellow color at room temperature indicates the presence of a very low concentration of the radicals. It is the presence of SO 3 F . radicals that gives S 2 0 6 F 2 the oxidizing ability that has been exploited in the synthesis of fluorosulfates in fluorosulfuric acid. The SO 3 F . radical has been isolated in an argon matrix [53], studied extensively by ESR [54], and its rotational and electronic spectra have been  14  reported [55]. The enthalpy of dissociation of S 2 06 F 2 has been reported to be between 91.3 and 97.6 k.l•mol -1 , depending on the method of determination. Bis(fluorosulfuryl) peroxide is non-ionizing, and has very poor solvating properties towards ionic solids. Many of the metals are often unreactive towards S20 6 F2 because of rapid passivation of the metal surface by a layer of insoluble metal fluorosulfate or fluoride or oxide. This obstacle may be overcome by using a solution of S 2 0 6 F 2 in fluorosulfuric acid. The solutions of S 2 0 6 F2 in HSO 3 F combine the oxidizing strength of S 2 06 F2 with the ionizing capacity of HSO 3 F and allow the extension of the working range of S 2 06 F 2 to +140°C, with a marginal attack on glass at the higher temperature. There are distinct advantages in using the HSO 3 F-S 2 0 6 F 2 solution, rather than S 2 06 F 2 alone, in metal oxidation reactions where the precursors are available in high purity, in the form of metal powders: a) the reactions are faster and generally go to completion; b) in addition to binary fluorosulfates, ternary fluorosulfates may be obtained easily in a one-step synthesis; c) isolation of solid products by filtration or removal of the liquid phase in vacuo is possible; and d) in the case of products that are soluble in fluorosulfuric acid, the removal of S 2 06 F 2 in  vacuo on account of its lower boiling point yields solutions that may be studied further.  THE FLUOROSULFATES The fluorosulfate anion is a weakly basic and a weakly coordinating ligand. It makes a suitable counteranion for high-valent, highly electrophilic metal and non-metal ions. The fluorosulfate radical may be viewed as a pseudohalogen, and the chemistry of S 2 06 F 2 displays many parallels to halogen chemistry, particularly to fluorine chemistry [56]. The synthetic routes to element fluorosulfates are often comparable to those for the halides. The principal features may be extended to similar anions of the type SO 3 X (X = CF 3 , Cl, etc.). There is a current overview on the synthetic aspects of fluorosulfate chemistry [57].  ^  15  Synthetic Routes to the Fluorosulfates There are four general approaches to element fluorosulfates: a) ligand substitution by solvolysis of corresponding metal derivatives in fluorosulfuric acid, HSO3F; b) metal oxidation or halide substitution with bis(fluorosulfuryl) peroxide (S20 6 F2) or the halogen fluorosulfates (XOSO2F, X = Cl or Br) in the presence, or absence, of HSO 3 F; c) insertion of SO 3 into the M-F bonds of metal fluorides; and d) miscellaneous methods such as the reaction of bis(fluorosulfuryl) oxide (S 2 05F 2) with alkoxides and the silver salt metathesis.  a) Solvolysis reactions in HSO 3 F: Of the above methods, solvolysis is almost exclusively the route of choice to the low valent 3d metal(II) fluorosulfates. Metal chlorides and carboxylates have been used as precursors, as shown below:  MC1 2 + 2 HSO 3 F ^ > M(SO 3 F) 2 + 2HC1^-36-  M(C 6 H 5 C00)2 + 4 HSO 3 F ^ M(SO3F) 2 + 2 C 6 H5COOH 2 + + 2 SO3F -^-37-  The volatility of HC1 and the solubility of the protonated carboxylic acid allow for the easy separation of the product. This method has been extended to metal sulfates, fluorides and acetates to yield the corresponding metal(II) fluorosulfates [58][59]. Goubeau and Milne have found that the reaction rate of the solvolysis of metal chlorides is enhanced by the presence of KSO 3 F [60]. The element fluorosulfate formed may undergo decomposition; the solvolysis of ZrC1 4 in HSO 3 F was reported to yield ZrF3SO3F [61]. The most obvious limitation of this method is the availability of suitable precursors. Some transition metal fluorosulfates obtained in this manner are listed in Table 1.2 below.  16 Table 1.2: Transition Metal Fluorosulfates obtained via solvolysis in HSO 3 F  Compound  Synthesis  Ref  Zr(SO 3 F) 4  Zr(CF3CO2)4 + HSO 3 F  [62]  V(SO 3 F) 3  VC13 + HSO 3 F + KC1  [63]  Mn(SO3F)2  Mn(CH3CO2)2 + HSO 3 F  Fe(S 0392  FeSO4 + HSO 3 F, 5 h  [64]  FeC1 2 + HSO 3 F, 1 d  [65][66]  [64][65]  Fe(SO 3 F) 3  FeC13 + HSO 3 F  [60]  Co(SO 3 F) 2  CoSO4 + HSO 3 F  [64]  Co[m-BrC 6 H 4 CO 2]2 + HSO 3 F Ni(SO 3 F) 2  NiSO4 + HSO 3 F Ni(CH 3 CO 2) 2 + HSO 3 F NiC1 2 + HSO 3 F, 14 d  Cu(SO 3 F) 2  CuSO4, Cu(CH 3 C00)2, CuC1 2 or CuF2 + HSO 3 F  [59][65] [64] [59][64] [65] [59][64]  Cu(C 6 H5CO 2 ) 2 + HSO 3 F  [65]  CuC1 2 + HSO 3 F  [65]  AgSO 3 F  Ag(CO2CF3) + HSO 3 F  [46]  Zn(SO3F)2  Zn(CH3CO2)2 + HSO 3 F  [64]  Zn(C 6 H5 CO 2 ) 2 + HSO 3 F  [65]  ZnC1 2 + HSO 3 F  [60]  Cd(C6H5CO2)2 + HSO 3 F  [65]  Cd(SO3F)2  17  b) Oxidation or ligand substitution by S 206 F 2 and XOSO 2 F (X = Cl or Br): The oxidation of metals by S 2 06 F 2 , in the presence or absence of HSO3F, has led to the synthesis of many binary fluorosulfates according to:  M + n/2 S20 6 F2  (HSO 3 F) ^>  M(SO 3 F),,^ -38-  The oxidation reaction is not limited to metals, and may also be extended to metal carbonyls. The reaction of bis(fluorosulfuryl) peroxide with metal carbonyls results in the partial oxidation of the displaced CO to CO2 by S 2 06 F2 (with S 2 05 F 2 formed as a byproduct) and requires an excess of the oxidizing reagent; it offers no advantage over the use of metal powders. Only three binary fluorosulfates have been obtained by the oxidation of metal carbonyls — Mn(SO 3 F) 3 [67], Cr(SO 3 F) 3 and Co(SO 3 F) 2 [68].  Mn2(CO) 1 0 + S2O 6 F2 Cr(CO) 6 + S2O 6 F2 Co2 (CO)8 + S 2 06 F2  60°C,20 d 6h  > Mn(SO3 F) 3  -39-  > Cr(SO 3 F)3  -40-  > Co(SO 3 F)2  -41-  An excess of S 2 O6F2 is also required in reactions with chlorides where the substituted chlorine is oxidised to C1SO 3 F and, sometimes, C10 2 S0 3 F. Chloride substitution is often incomplete, resulting in the formation of compounds like TiC1 2 (SO 3 F) 2 [69]. The use of metal oxides or carbonates as substrates is not recommended, due to their propensity to form metal oxyfluorosulfates [70]. The binary fluorosulfates of the transition metals obtained via oxidation with S 2 O6 F 2 are listed in Table 1.3. In addition to the compounds mentioned in Table 1.3, the pentavalent fluorosulfates of Nb and Ta, Nb(SO 3 F) 5 and Ta(SO 3 F) 5 have been generated in situ, but were not isolable [71].  18  Table 1.3: Transition Metal Fluorosulfates obtained  via oxidation with S 2 06 F 2 in the presence  or absence of HSO 3 F  Compound  Synthesis  Ref  VO(SO 3 F) 3  V(CO)6 + S 2 0 6 F 2  [68]  NbO(SO 3 F) 3  NbC15 + S 2 06 F2  [72]  NbF 2 (SO 3 F) 3  Nb + S 2 06 F 2 in HSO 3 F  [71]  TaF 4 (SO 3 F)  Ta + TaF 5 + S 2 06 F 2 in HSO 3 F  [71]  TaO(SO 3 F) 3  TaC15 + S206F2  [72]  Cr(SO 3 F) 3  Cr(CO)6 + S 2 06 F2 at 60°C, 20 d  [68]  Mo0 2 (SO 3 F) 2  Mo(CO)6 + S 2 0 6 F 2  [73]  Mn(SO 3 F) 3  Mn + S206F2  [74]  Mn(SO 3 F) 2 + S 2 0 6 F 2  [74]  Mn(SO 3 F) 4 a  Mn2(CO)10 + S 2 06 F 2  [68]  MnO(SO 3 F)  MnCO3 + S 2 0 6 F 2  [70]  Re0 2 (SO 3 F) 3  Re + S 2 06 F 2  [72]  Re0 3 (SO 3 F)  Re + S 2 0 6 F 2  [72]  Ru(SO 3 F) 3  Ru + S 2 0 6 F 2 in HSO 3 F, 60°C, 24 h  [75]  Os(SO 3 F) 3  Os + S 2 0 6 F 2 in HSO 3 F, 60°C, 3 d  [76]  WO(SO3F)4  W(CO)6 + S 2 0 6 F 2  [77]  WC1 6 + S 2 O 6 F 2  [77]  Co(SO3F)2  Co2(CO)8 + S2O 6 F2 , 6 h  [68]  Rh(SO 3 F) 3  Rh + S 2 0 6 F 2 in HSO3F, 130°C, 21 d  [76]  Ir(SO 3 F) 3  Ir + S 2 0 6 F2 in HSO 3 F, 140°C, 6 d  [78]  Ir(SO 3 F) 4  Ir(SO3F)3 + S 2 0 6 F 2 in HSO 3 F, 60°C, 6 h  [78]  NiO(SO3F)  NiCO3 + S 2 0 6 F 2  [70]  19  Table 1.3: Transition Metal Fluorosulfates obtained via oxidation with S 2 0 6 F2 in the presence or absence of HSO 3 F (...cont'd)  Compound  Synthesis  Ref  Pd[Pd(SO3F)6]  Pd + S 2 06 F2 in HSO3 F, 80°C, 24 h  [79]  Pt(SO3F)4  Pt + S 2 06 F 2 in HSO3F, 120°C, 2 d  [80]  Ag(SO3F)2  Ag + S 20 6 F 2 at 70°C, 7 d  [81]  Ag + S 2 0 6 F 2 in HSO 3 F, 25°C, 30 m  [46]  Ag 2 O, Ag2CO3, AgSO 3 F, AgSO 3 CF 3 + S 2 06 F2  [46]  Ag20(S0392  Ag2CO3 or Ag 2 O + S 2 06 F 2  [70]  Au(SO 3 F) 3  Au + S 2 06 F2 in HSO3F, 3 h  [82][83]  Zn(SO 3 F) 2  Zn + S 2 06 F2 in HSO 3 F, 90°C, 21 d  [84]  Cd(SO 3 F) 2  Cd + S 2 0 6 F 2 in HSO 3 F, 90°C, 28 d  [84]  Hg(SO 3 F) 2  Hg + S 2 06 F 2  [65][85]  Hg + S 2 06 F2 in HSO 3 F  [84]  'This compound was later shown to be Mn(SO 3 F) 3 [74].  In addition to bis(fluorosulfuryl) peroxide, the halogen fluorosulfates, in particular BrSO 3 F and C1SO 3 F, have also been used to obtain some fluorosulfates from the elements or their chlorides and oxides (See Table 1.4). The use of BrSO 3 F offers no advantages over the use of S 2 0 6 F 2 because: a) bis(fluorosulfuryl) peroxide is needed to produce bromine fluorosulfate; b) longer reaction times are needed when BrSO 3 F is used as a reagent; and c) when BrSO 3 F is used very often intermediates are formed which have to be pyrolyzed in order to yield the respective binary fluorosulfate derivatives. This method, however, has been used in the successful synthesis of the fluorosulfates of Ag [46], Au and Pt [86] and Pd [79] by direct oxidation of the metals with BrSO 3 F.  20  The two methods discussed above, solvolysis and oxidation, are the most common routes to binary fluorosulfates, and may be viewed as being complementary to one another. While the former has seen much use in the synthesis of the lower valent binary fluorosulfates on account of the ready availability of the necessary precursors, the latter has been extensively exploited in the synthesis of the higher valent fluorosulfates due to the strong oxidizing ability of S 2 06 F2 . Higher valent binary fluorosulfates may also be reduced by careful thermal decomposition to lower valent compounds with the elimination of either S 20 6 F 2 , or S 20 5 F 2 and 0 2 . This method has been exploited in the synthesis of Au(SO 3 F) 2 as described in Appendix A.  c) Insertion of SO3: The insertion of SO 3 into M-F bonds has been useful in the synthesis of the ionic fluorosulfates of the alkali metals and alkaline earth metals [22][87][88]. Although this method has been used to synthesize Cr(SO 3 F) 3 , the starting material used was CrF 5 and not CrF 3 as one might expect [89].  CrF 5 + 5 SO 3 ^ > Cr(SO 3 F) 3 + S 2 06 F 2^-42-  However, this method is not easily applicable to the preparation of higher valent fluorosulfates because incomplete substitution is often encountered resulting in fluoro-fluorosulfato derivatives of the type MF n (SO3F) m .  MF 5 + n SO 3  > MF5_n(SO3F)n ^ -43-  The reaction of SO 3 on UF 6 at -60 to -50°C in CFC1 3 gave, along with S 2 06 F2 as a byproduct, a material of the composition UF(SO 3 F) 4 in which the uranium is in the +5 oxidation state [90]; at higher temperatures, however, UF 2 (SO 3 F) 3 is reportedly formed [91].  21  d) Miscellaneous methods: The reaction of S 2 05 F 2 with Ti(OCH 3 ) 4 has resulted only in disubstitution [92] and is the only reported attempt of its use in the synthesis of transition metal fluorosulfates. Bis(fluorosulfuryl) oxide, S 2 05 F2 , has been used successfully in the preparation of Bi(SO 3 F) 3 [92]. In a few cases, the silver salt metathesis from metal halides has been exploited with some success [93] in preparing C104, BF4 and PF6 derivatives, the formation of the silver halide being the driving force behind the reaction. However, there are some disadvantages in preparing fluorosulfate derivatives according to:  MX /1  + n AgSO 3 F  > M(SO3F) n  + n AgX^-44-  AgSO 3 F is difficult to make, multiple substitutions (n > 2) are difficult to carry out, and product separation is an arduous process. However, this method has been successfully used in:  25°C M(CO) 5 X + AgSO 3 F ^ > M(CO) 5 S0 3 F + AgX^-45CH 2 C1 2 M = Mn, Re; X = Cl, Br  where the reaction proceeds quickly for Re but slowly for Mn [74]. Product separation from AgX is possible because the M(CO) 5 S0 3 F compounds are soluble in CH 2 C1 2 . Brazier and Woolf [94] have studied the effect of boiling HSO 3 F on various transition metals, which produce SO 2 and, on occasion, even elemental sulfur as a byproduct. This merits attention only for the synthesis of Cu(SO 3 F) 2 and possibly AgSO 3 F, either because HSO 3 F often cannot oxidize the metal or because an insoluble layer of the fluorosulfate passivates the metal surface. The transition metal derivatives obtained via routes involving BrSO 3 F oxidation, SO 3 insertion, and the controlled pyrolysis of the higher valent fluorosulfates are listed in Table 1.4.  22  Table 1.4: Miscellaneous Routes to the Transition Metal Fluorosulfates  Compound  Synthesis  Ref  Synthesis by SO 3 insertion into M-F bonds. NbF 3 (SO 3 F) 2  NbF5 + SO 3  [95]  TaF3(SO3F)2  TaF5 + SO3  [95]  Cr(SO 3 F) 3  CrF5 + SO3  [89]  WF 2 (S 03F)4  WF6 +  [96]  Ag(SO3F) 2  AgF2 + SO3  [46]  Ag + BrF 3 + SO 3  [97]  Ti(OCH3)4 + S 2 O5F2 at 60°C, 2 d  [92]  SO 3 from M/A: WF 6 4.5(S0 3 )  Addition of S 2 0 5 F 2 . Ti(OCH3) 2 (SO3F)2  Synthesis by pyrolysis or metal powder reduction. Ir(SO 3 F) 3 ,  Ir(SO3F)4 at 120°C  [78]  Pd(SO 3 F) 2  Pd[Pd(IV)(SO3F)6] at 160°C  [98]  AgSO 3 F  Ag(SO3F)2 at 210 ° C  [81]  Au(SO 3 F) 2  Au(SO3F)3 + Au in HSO 3 F, 22 d  [99]  Reactions with BrSO 3 F Pd(SO3F)2  Pd + BrSO 3 F, 110°C, 14 d  Pt(SO3F)4  Pt + BrSO 3 F, 95°C, 21 d  [86]  Ag3(SO3F)4  Ag, Ag 2 O or AgSO 3 F + BrSO 3 F 25-70°C, 2 d  [46]  Ag(SO 3 F) 2  Ag, Ag 2 O, or AgSO 3 F + BrSO 3 F  [46]  Au(SO 3 F) 3  Au + BrSO 3 F, 65°C, 1 d  [86]  [79][100_  23  Decomposition of the Fluorosulfates The decomposition of binary fluorosulfates proceeds in four general ways: a) elimination of S 2 05 F2 resulting in oxyfluorosulfates or oxides (eg. V, Nb, Ta, Mo and Re); b) elimination of sulfur trioxide, SO 3 , producing fluoride-fluorosulfates or fluorides (eg. Nb and Ta); c) elimination of SO 2F 2 to give metal sulfates (eg. Ba and Sr); and d) reductive elimination of SO 3 F . radicals to form S 2 06 F 2 and a lower valent binary fluorosulfate (eg. Au, Pd and Ag). Of the above, a) and b) are the most common decomposition modes; both may occur simultaneously. The low valent fluorosulfate derivatives of the noble metals have been obtained by the controlled pyrolysis of the corresponding higher valent derivatives.  STRUCTURAL ANALYSIS OF THE FLUOROSULFATES The fluorosulfate ion, like the fluoride ion, forms ionic solids with electropositive metals. The fluorosulfate group, like the fluoro group, can form polar covalent bonds (eg. F2N-SO 3 F vs. F 2 N-F), or can act as a coordinated ligand in a mono-, bi-, or tridentate fashion. With respect to structural elucidation, two principal differences arise between the fluorides and fluorosulfates: a) the fluoride ligand is monoatomic, similar to Cl - or other halides, and much use has been made of X-ray diffraction methods (both on single crystals and powders). This is not so for the fluorosulfate ligand, where only a few single crystal structures are available, and powder diffraction is not used to a significant extent because of the difficulty in inferring structural information; and b) the fluorosulfate group is a molecular (polyatomic) group and vibrational spectroscopy becomes very useful because the vibrational frequencies extend well into the mid infrared region and permit the use of group frequency assignments which, in turn, reveal coordination modes. The vibrational frequency range of metal fluoride modes is far narrower; stretching modes occur well below 700 cm -1 , which make detection and interpretation difficult, particularly in solids, where the lattice vibrations manifest themselves. In addition, the SO 3 F group is a good Raman scatterer on account of strong covalent bonding  24  within the group, whereas M-F vibrations are strongly polarized and frequently poor Raman scatterers. Many fluorosulfates are insoluble in HSO3F and many other common solvents, indicative of a high degree of polymerization. This, in combination with the high reactivity of the fluorosulfates, has frequently precluded their characterization by single crystal X-ray diffraction studies. MOssbauer spectroscopy has been used to a limited extent to characterize fluorosulfate derivatives containing Fe or Sn. NMR spectroscopy, which has been used extensively for studying liquids and solutions, is more recently and to a lesser extent applied in the study of solids. Vibrational spectroscopy has been the most widespread technique in the investigation of fluorosulfates. The samples may be in the solid, liquid or gas phase. The use of silver halide cells and moisture-free conditions has facilitated the recording of the infrared spectra of many fluorosulfates. The advent of FT-IR spectrometers has been helpful in obtaining the spectra of more reactive or opaque species. Raman spectra can, however, be obtained on solid or liquid samples sealed in glass tubes. When intense color, limited thermal stability and fluorescence are not encountered, one may obtain relatively simple Raman spectra since only the fundamental vibrations are observed; overtones and combination bands are not usually seen. The fluorosulfate group is known to exhibit several bonding or coordination modes. The effects of coordination have a direct effect on the number of fundamentals and their frequencies in the infrared region. A brief look at some of the previously reported structures serves as a useful preamble to their characterization by vibrational spectroscopy.  Structures of some fluorosulfates Although there is a large number of inorganic fluorosulfate derivatives known and reported, there is a paucity of single-crystal X-ray diffraction studies. The few structures that are available, however, provide good examples of the various bonding or coordination modes of the fluorosulfate group, and some of the representative structures are discussed below.  25  The ionic fluorosulfate group is exhibited in the crystal structure of KSO 3 F [101], wherein all the S-0 bond lengths and O-S-0 angles are the same at 1.43(1) A and 112.9(7)°, respectively. The S-F bond is 1.58(2)  A and the 0-S-F angles are all 105.8(7)°.  The structure of gold tris(fluorosulfate) [102] is the first example to confirm the presence of both terminal monodentate and symmetrically bridging bidentate fluorosulfate groups in the same molecule. The compound exists as a centrosymmetric dimer in the solid state (See Fig. 1-1), and the 8-membered ring formed by the two bridging fluorosulfates and the gold centres is in a chair conformation. Both the Au(III) centres are in a slightly distorted square planar environment, brought about by the terminal monodentate groups being bonded more strongly to the Au centres (average Au-0 = 1.957(16) A) than the bridging bidentate groups (average Au-0 = 2.018(13) A). In addition to these four Au-0 bonds, there are two longer Au•••0 contacts (2.960(7) and 2.757(8) A) from other dimers above and below the Au0 4 square plane. These "axial" oxygen atoms exhibit significant departure from linearity (0-Au-0 = 158.6(3)°), resulting in a distorted 4-2 arrangement about the Au(III) centres. There is significant thermal motion of the non-bonded oxygens and fluorine in the terminally bonded monodentate fluorosulfate groups. For both types of fluorosulfate groups, the S-O b bonds are longer than the S-O t bonds, and this is also reflected in the vS-0 in the vibrational spectra (vide infra). The recently obtained crystal structure on SbF(SO 3 F) 2 [103] gives an example of an asymmetric bidentate bridging fluorosulfate group. Each fluorosulfate group is attached to the antimony atoms by one long bond (ay. 2.53 A) and one short bond (ay. 2.12 A). This type of bonding is also reflected in the S-0 bonds; there are three S-0 bond lengths seen in the fluorosulfate group: 1.41  A, 1.436 A, and 1.48 A. The longest S-0 bonds correspond to the  shortest Sb-0 bonds (2.122(8) A), the intermediate S-0 bonds correspond to the longer SbO bonds (2.52(5) A), and the shortest S-0 bonds (1.407(2) A) with longer intermolecular secondary contacts between the Sb and 0 (3.03 A).  26  Fig. 1-1. Crystal structure of gold tris(fluososulfate), [Au(SO 3 F) 3 ] 2 (Ref. 102)  27  The crystal structure of (CH 3 ) 2 Sn(SO 3 F) 2 also exhibits bidentate bridging fluorosulfate group [104]. The molecules form a sheet polymer in which the sheets are held together by weak van der Waal's forces, and the general arrangement is akin to that in (CH 3 ) 2 SnF 2 . The tin atom may be viewed as being octahedrally coordinated with Sn-C distances of 2.08(1)  A.  For the SO 3 F group, the three S-0 bond distances are equal (1.47(1) A) within experimental error, the S-F bond length is 1.56(1)  A; the 0-S-F angles are 107° and the O-S-0 angles are  112°. The minimal departure of the SO 3 F group from C3 v symmetry, the long Sn-0 distances (2.24(1) A)and the short Sn-C distances allow interpretation of the structure in terms of (CH 3 ) 2 Sn 2+ cations and SO3F - anions where bridging fluorosulfates connect the tin atoms via weak Sn-O bonds. The crystal structure of Sn(SO 3 F) 2 [105] displays a 3-dimensional network, with the fluorosulfate groups acting as tridentate covalently bridging ligands. The primary geometry about the Sn centre is S4, which is similar to other oxygen-coordinated Sn(II) compounds [106][107]. In addition to these four oxygen atoms, there are two longer contacts (ay. 2.7 A) with oxygen atoms on the fluorosulfate group. The three S-0 bonds are not equal and, like in the case of SbF(SO 3 F) 2 , the shorter S-0 bonds correspond to the longer Sn-0 bonds. The fluorosulfate groups show significant departure from C3 v symmetry; this had been predicted by spectroscopic results [108][109]. We have recently obtained crystal structures on some carbonyl derivatives of transition metal fluorosulfates. The structure of cyc/o-[Pd2(A-00)2](S03F)2 exhibits the presence of ionic bidentate bridging fluorosulfates and that of cis-Pd(C0) 2 (SO 3 F) 2 shows the presence of covalent monodentate fluorosulfate groups. Hence, the few crystal structures that have been obtained provide examples for most of the common coordination modes of the fluorosulfate group, and vibrational data can be used, with some confidence, to postulate the structures of other fluorosulfate derivatives for which crystal structures are not available.  28  Vibrational analysis of the fluorosulfate group  The analysis of the vibrational spectra of the fluorosulfate group provides important information regarding the type of coordination and the extent of interaction. There are distinct advantages to working with the polyatomic fluorosulfate group as compared to the monoatomic F - and Cl ligands: -  i)  the fundamental modes span the mid- to far-infrared region (1500 - 300 cm -1 ).  ii)  the stretching modes used for structural characterization occur at 1500-700 cm -1 , distinguishable from M-0 stretches.  iii)^the S-0 and S-F bonds are good Raman scattering groups and allow complementary information to be obtained. The molecular structures of the fluorosulfates reported so far show ionic, monodentate, bidentate (symmetrically or asymetrically) bridging or tridentate bridging fluorosulfate groups. It is assumed that the polydentate fluorosulfate groups are bridging rather than chelating, and there is, up to the present time, no evidence for a chelating SO 3 F group. The general physical and spectroscopic properties of these compounds are consistent with either oligomeric structures such as that for Au(SO 3 F) 3 [102] or polymeric structures like that of (CH 3 ) 2 Sn(SO 3 F) 2 [104]. The C3, symmetry of the free SO 3 F" ion can be reduced to C s or even C 1 upon coordination. When one or two of the oxygens on SO 3 F are coordinated, this results in an increase in the number of vibrational modes from 6 (3A 1 + 3E) to 9 (6A' + 3A"). Further coordination once again increases the symmetry to C3, and reduces the number of vibrational modes to six. The presence of strongly polarizing or non-spherical counter cations can partially lift the degeneracy of the three E modes in an ionic SO3F group resulting in further band proliferation. This also occurs due to site symmetry effects which are seen when the SO 3 F - ion is in a crystallographic site having a symmetry lower than C 3 ,,. All the vibrational modes of the SO 3 F group are both infrared and Raman active.  29  In the vibrational analysis of the fluorosulfates, we concentrate on the stretching modes more than on the bending modes because the stretching modes occur over a wider range in a region of the spectrum that is usually not obscured by metal ligand bands. In addition, the stretching modes are more dependent on the bonding of the fluorosulfate group than the bending modes, and are therefore more diagnostic. Within the two groups of spectra (Fig. 1-2) the differences in band positions due to varying bond strengths are monitored in the "diagnostic bands". Coordination of the oxygen atom in the fluorosulfate group weakens the S-0 bond and lowers the vS-O. When electron withdrawal occurs, the uncoordinated S-0 and S-F bonds are strengthened, thus raising the vS-O and vS-F. The shaded bars in Fig. 1-2 indicate the diagnostic vS-O and vS-F which occur in unobscured regions of the spectrum and give a useful indication of the coordination mode of the fluorosulfate group.  Ionic and ionic perturbed: Purely ionic salts such as alkali metal fluorosulfates MSO 3 F  (M = K [101](110], Na and Cs [110]) exhibit spectra that are in good agreement with a normal coordinate analysis of the SO 3 F group in the A l ground state [111]. The observed six band spectra (3A 1 + 3E) are relatively simple to interpret, with the low vS-F band  (ca. 750 cm -1 ) being the most diagnostic band. The presence of non-spherical cations such as NO or strongly polarizing cations such as Li + , lifts the degeneracy of the three doubly degenerate E modes resulting in band proliferation as seen in the spectra of NOSO 3 F [112] and LiSO 3 F [110]. These spectra are distinguished by the presence of a closely spaced pair of v as S0 2 and vsS0 2 and a low vS-F (< 800 cm -1 ).  Covalent tridentate: The covalent tridentate bridging fluorosulfate group is commonly seen in  polymeric compounds of the type Pd(SO 3 F) 2 and Zn(SO 3 F) 2 , which appear to be related to the CdC1 2 layered structure. This coordination mode is differentiated from that of the ionic  $: E v)  Bonding or coordination modes  Frequency Range (cm') stretching bands  deformation bands  v. (S-0)^ v (S-F) 1^1^1^I^r^A v„ (S-0)  covalent tridentate  i^i^p S. (SO) (SO,F) f^I  I  1  5„ (SO)  U v. (S-0)^v„ (S-0)^V (S-F) CI^i^I^ •^A  purely ionic  ionic perturbed  , ,., '  covalent monodentate  v. (SO 2) =1^v (S-0)^v (S-F) 1 ^ i^ = Km Vsy. (SO2) v (S-0) v. (SO2)^vv^ „, (SO2) I ^ I ^ I ^I I^I  v (S-F)  covalent bidentate  V. (SO2)^ V (S-0)^ V (S-F) f-------1^ I^ 1 2=21 V sy. (SO2)  I  5. (SO)^p (SO,F) I^IN^I^1 5, (SO,) 5 (SO2)  1^Y  (OSF)^y (SO,F) o^I^i^ NI^t (SO 2F) Y (SO2) y (SO 2F) 5 (SO2)^y (OSF)^= = i^I^1^I 1 Y (S02) t (SO2F)  5 (SO)^y (OSF)^t (SO 2F) = =^^^E^II^1  Y (SO2)^y (SO,F)  ^^^^ 1 ^ I ^ I ^ I ^  1400  1200  Fig. 1-2. Vibrational frequencies of the fluorosulfate group  1000  800  600  ^  400  31  fluorosulfate group by the position of the vS-F which occurs well above 800 cm -1 . The S-0 bands are seen in a narrower range than for the ionic fluorosulfates. In addition, the band positions of the SO3 stretching modes are lower. The C3 v symmetry of the anion is retained unless Jahn-Teller distortion occurs as in the case of Cu(SO 3 F) 2 [59][60][65].  Covalent bidentate: There are nine fundamentals for the covalent bidentate SO3F group. The vS-O is at higher frequency, and the symmetric and asymmetric vSO2 are found close together between 1150 and 1200 cm -1 . Bidentate bridging groups are found in binary fluorosulfates of the type M(SO3F)3 (M = Ge, Fe and Mn), and ternary compounds like (CH3)2Sn(SO 3 F)2 and MF 2 (SO3 F) 2 (M = Sn and Ge).  Covalent monodentate: This type of bonding is commonly seen in the halogen fluorosulfates XSO3F (X = F, Cl, Br), and in the hexafluorosulfatometallate anions present in ternary salts such as Cs2 [M Iv (SO3F)6]. The presence of two vSO2 bands in the 1000-1450 cm -1 region along with a characteristic vS-O below ca. 1000 cm -1 and a vS-F at ca. 850 cm -1 is typical of this type of SO3F group. The only reported tetradentate fluorosulfate group is postulated for Ti 3 Cl io (SO 3 F) 2 [69] based on the presence of an extremely low vS-F (660 cm -I ) attributed to a weakened S-F bond resulting from coordination of the fluorine to the metal center. This is the lowest value for vS-F reported so far.  SUMMARY To summarize, the fluorosulfate group exhibits a strong similarity to fluorine that is also reflected in their chemistry. However, there are some contrasts that should be mentioned at this juncture: a) Higher oxidation states, and very occasionally higher coordination numbers, are observed for the fluorides than for the fluorosulfates in binary transition metal complexes; b) Transition metal fluorides exhibit a higher thermal stability than their fluorosulfate  32  counterparts (evidenced in the high temperature syntheses of the fluorides); c) With the inherent difficulty in obtaining single crystals, vibrational spectroscopy plays an important role in the structural characterization of fluorosulfates; and d) The reagents used in the synthesis of fluorosulfates, while analogous to those used in the synthesis of fluorides, are less corrosive and, unlike the fluoride reagents which require fluorinated polymers or metal equipment, can be handled safely in glass apparatus.  33  REFERENCES  1. S. Arrhenius, Z. Phys. Chem., 1 (1887) 631. 2. J.N. BrOnsted, Red. Tray. Chim. Pays Bas., 42 (1923) 719. -  3. H.P. 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Alleyne, K. O'Sullivan and R.C. Thompson, Can. J. Chem., 52 (1974) 336 110. A. Ruoff, J.B. Milne, G. Kaufmann and M. Leroy, Z. Anorg. Aug. l Chem., 372 (1969) 119. 111. S.P. So, Mol. Phys., 23 (1972) 1147. 112. A.M. Qureshi, H. Carter and F. Aubke, Can. J. Chem., 49 (1971) 35.  38  CHAPTER 2. EXPERIMENTAL  General experimental techniques, instrumentation and the sources of the common starting materials will be described in the following sections. The majority of the compounds involved in this work are extremely moisture sensitive, and have to be handled appropriately in an atmosphere free of residual moisture. This involves the use of an inert atmosphere box for the manipulation of solids and non-volatile liquids, standard vacuum-line techniques for handling volatile materials and, in addition, the development of more specialized, innovative methods and the design of appropriate equipment. The last two items will be dealt with in greater detail in this chapter.  GENERAL EQUIPMENT Utility Vacuum Lines and Glassware  All ground glass connections were lubricated with Fluorolube grease series 25-10M, CF 2 C1(CF 2 CFC1) n CF2C1 (Halocarbon Products Corporation). Its low volatility and reactivity towards the chemicals used in this field make it more suitable than hydrocarbon grease (Apiezon). Another suitable alternative may be found in Krytox (duPont, Canada). All the stopcocks used were greaseless Teflon stem valves obtained from Kontes (New Jersey, USA) or Young (London, UK). Teflon coated stir bars were used in conjunction with various magnetic stirrers to enable mixing of the reagents. A standard 90 mm manifold Pyrex vacuum line of known internal volume equipped with stopcocks and B10 ground glass joints and fitted with a Setra pressure transducer model 280E (Setra Systems Inc., Acton MA, USA) was employed for the manipulation of volatile reagents and reaction products. The trap-to-trap purification of products was accomplished on the vacuum line shown in Fig. 2-1 with three glass traps (Fig. 2-1) used in series so that each one of them could be bypassed at any given time. Vapor density measurements were used to determine the molecular weights of volatile reaction products.  39  ===pneflPafigerOf 5.7,W • AMMA•32.3:W %OM a A s ""  B14  B14 Shaded area = 50 mL \ _1  27 cm  5 cm  Fig. 2-1. Vacuum line and trap used for manipulation of liquids and volatile materials.  40  Reactions were carried out in either tubular or round-bottomed one-part Pyrex reactors of varying size fitted with Teflon stem stopcocks (Fig. 2-2). Reactions involving unpredictable reactants or high temperatures and internal pressures were performed in thickwalled (3 mm) glass reactors. For most of the other reactions, standard wall (2 mm) reactors were used. High temperature reactions over a long period of time were carried out in sealed ampoules of various sizes. In addition, reactive volatile reagents were flame-sealed in glass ampoules and stored at liquid nitrogen temperature.  Fig. 2-2. Reaction vessels used in synthesis.  41  Ampoule Key Volatile samples were sealed in glass ampoules of varying dimensions fitted with 6.0 mm o.d. Pyrex necks. The use of an ampoule key (Fig. 2-3a) enabled the sealing and re-opening of these ampoules under vacuum as described before [1].  Sublimation Apparatus A sublimation apparatus for vacuum sublimation as shown in Fig. 2-3b was used to purify airsensitive materials by vacuum sublimation. The apparatus may be opened under nitrogen or argon inside an inert atmosphere box to recover the material.  Vacuum Filtration Apparatus An elaborate, if not awkward, piece of apparatus (Fig. 2-4) was used for the isolation of insoluble, moisture-sensitive products from the solvents and soluble impurities. The design here enables the rotation of the whole apparatus about the B14 cone allowing the transfer of the solution through the glass frit. Sequential manipulation of the stopcocks enables the introduction of nitrogen into the upper section, followed by its evacuation through the glass frit  via the lower chamber. The inherent advantages of this design over previous ones are its relative simplicity and strength.  Metal Vacuum Line and Vessels Reactions involving anhydrous HF were carried out on a 0.25" o.d. Monel vacuum line equipped with a pressure transducer attached to a micron gauge. The line was equipped with a Monel bridge to enable vacuum transfer of HF in and out of the reactor. Flexible extensions to this line were made using Monel or copper tubing with either brass or stainless steel Swagelok fittings. The line was passivated at high temperature and/or high pressures of fluorine prior to use. The reactors used were of two types. The first involved a 6" long Kel-F cylinder machined out of a 1.5" o.d. Kel-F rod (Argonne National Lab) that was seated onto a  42  H 2 O in 4-B10 Joint H 2 O out  01  --ON.  4mm Kontes Teflon Stopcock B24 or B45 Joint  Vacuum  B45 Joint  Stressor BI4  B 14 taper a)^  Viton 0-ring  b)  wosher  Fig. 2-3. a) Ampoule key used for opening and resealing glass vials under vacuum. b) Sublimation apparatus used for purification of air-sensitive compounds.  43  50-100mL Pyrex flask  — B19 joint  Glass frit (medium or coarse porosity)  Kontes 4mm Teflon stem stopcocks  Kontes 4mm Teflon stem stopcocks  © B19 joint  50-100mL Pyrex flask  Fig. 2-4. Filter apparatus for filtration of air-sensitive samples.  44  one-piece Monel adapter by means of a copper ferrule held in place by a brass lock-nut. The Monel adapter was welded onto a Whitey IRS4 stainless steel valve which provided an attachment to the vacuum line (Fig. 2-5). The second reactor was a one piece Teflon reactor, machined from a single Teflon block, attached to a stainless steel plate equipped with a Whitey IRS4 stainless steel valve by means of a stainless steel flange.  Inert Atmosphere Box All non-volatile, air-sensitive reaction materials and products were handled inside a Vacuum Atmospheres Corp. Drybox Model DL-001-S-G Dri-Lab equipped with a Model HE-493 DriTrain filled with molecular sieves. The drybox was filled with "L Grade" nitrogen (Linde, Union Carbide) and was periodically regenerated by heating the sieves and copper catalyst in the Dri-Train and passing a mixture of 10% H2 in N2 (Linde, Union Carbide), after which the pump oil was replenished. All reactions were monitored by weight differences in the reaction vessels. A Mettler Gramatic analytical balance #1-910 (max. 200 g; ± 0.2 mg) was used for the general weighing of reagents and products. Heavier stainless steel reactors were weighed on a toploading Sartorius Type 1104 (max. 1000 g; ± 0.5 g). A top-loading Mettler PC440 balance was used to measure out samples inside the drybox.  INSTRUMENTATION Infrared Spectrometers Depending on the sample and cell windows used, the infrared spectra were recorded on a variety of spectrophotometers. A Perkin-Elmer Model 598 grating IR spectrophotometer was used for solid or gaseous samples when AgBr or AgCI windows were used. A Bomem Michelson MB102 FT-IR was interfaced to an IBM compatible 80386 computer equipped with an Intel 80387 floating point coprocessor using a proprietory 16-bit multiplexed data  45  Whitey Valve (1RS4-316)  Monel Alloy Top  ))  Kel-F Tube  6"  iv  Copper Ferrule  Brass Locknut  Fig. 2-5. Kel-F reaction vessel for reactions in HF.  1\  ^  1  A  r  46  acquisition board. Galactic LabCalc 2.24 (Galactic Industries, NH, USA) and BomemEasy 1.45 (Bomem, Quebec, Canada) software packages were used to acquire and analyze the data. On occasion, a Nicolet MX-1 spectrometer operating between 4000 and 400 cm -1 and a Bruker IFS-66v instrument with a range of 4000-50 cm -1 were also employed. For all infrared spectroscopic measurements, the spectral accuracy was 4 cm 1, unless otherwise mentioned; the spectral resolution was between 1 and 4 cm -1 . Solid samples were pressed as thin films between AgBr or AgC1 windows (Harshaw Chemicals, USA) or IR grade silicon windows (Wacker Chemie, Burghausen, Germany). The high reactivity of most fluorosulfates prevents the use of the commonly used alkali halide windows (NaCI, KC1 or CsI) or any mulling agents such as Nujol, hexafluorobenzene or hexafluorobutadiene. Occasionally, Fluorolube oil was used as a dispersant when the compounds exhibited limited reactivity towards it. Gaseous samples were contained in a 10 cm glass gas cell fitted with 1 mm thick AgBr windows, cut out from rolled sheets of the silver halide (Harshaw Chemicals, USA), or IR grade silicon windows (Wacker Chemie, Burghausen, Germany). The gas cell was equipped with a cold finger to enable the condensation of all volatile materials at any given temperature into the cell (see Fig. 2-6). This is extremely useful when reactions are carried out on a small scale.  Fig. 2-6. Gas cell for obtaining infrared spectra on gaseous samples.  47  Raman Spectrometers Raman spectra were obtained on a Spex Ramalog 5 equipped with an argon ion laser (Spectra Physics 164) with the liquid samples contained in 5 mm o.d. NMR tubes or 6 mm o.d. glass tubes. Solid samples were sealed in 2.0 mm o.d. glass capillaries. The 514.5 nm line was used for excitation. Some spectra were obtained on a Bruker FT-Raman Instrument type FRA 106 equipped with a Nd-YAG laser (Prof. Dr. Schnoekel, Munich; Prof. Dr. Schraeder, Essen; Prof. Dr. Homborg, Kiel; or Dr. Sawatzki, Bruker Instruments, Karlsruhe).  Nuclear Magnetic Resonance Spectrometers 1H  and 19 F NMR spectra were obtained on a Varian XL-300 spectrometer operating at 300  and 282 MHz, respectively, fitted with a 5mm bore multinuclear switchable probe, or on a Bruker CXP-200 automated spectrometer operating at 200 and 188 MHz, respectively. The spectra were recorded unlocked on neat liquid samples, with CFC1 3 in d 6 -acetone used as an external reference for the 19 F chemical shift. A 5 mm bore broad band probe with a tunable range of 30-122 MHz was used on the XL-300 to obtain the 77 Se NMR spectra with the instrument operating at 57 MHz. A solution of SeO 2 in D 2 0 was used as an external reference. The  95 Mo  spectra were obtained on a  Bruker WH-400, fitted with a 10.0 mm broadband probe, operating at 26 MHz. Spectra were referenced to a solution of aqueous molybdate in D 2 0. Samples were held in 5.0 mm o.d. NMR tubes either sealed off or equipped with rotationally symmetric 5.0 mm Teflon stem valves (Young, London; V-NMR /5). Due to the absence of a deuterated solvent, the magnet was shimmed to a tube containing d 6 -acetone. All chemical shifts at lower field to the references are quoted as positive values.  Electron Spin Resonance Spectrometer The ESR spectra were obtained on an X-band homodyne spectrometer equipped with a Varian 12-inch magnet and a MkII Fieldial control. The samples were held at -170°C using a Varian  48  E-257 temperature controller. An Ithaco Dynatrac 391A lock-in amplifier was used to obtain phase-sensitive detection at 100 kHz. The microwave frequency was measured with a HP5246L frequency counter equipped with a HP5255A plug-in. Field calibration was accomplished using a Varian Gaussmeter, the output of which was collected by an IBM XT computer. Data aquisition was carried out using a Qua-Tech 12-bit data aquisition board (ADM 12-10) together with a Qua-Tech parallel expansion board (PXB-721) incorporated into an IBM XT computer. The relevant data processing software is as described previously [2][3]. The absolute field was corrected for the placement of the gaussmeter probe by calibrating against peroxylamine sulphate in aqueous solution and is in error by ca. 0.01 G. The precision of the data is typically ca. 0.02 G for 2K points over a 100 G sweep which, together with the ca. 10 kHz error in the microwave frequency, corresponds to an error of ca. 0.00002 in the g-values quoted. The simulations were performed using FORTRAN programmes written by Dr. P.S. Phillips (Department of Chemistry, UBC; now Okanagan University College, BC) based on the equations of R.E.D. McLung [4] and EPR analysis software (POWD and QPOW) supplied by the Illinois ESR Research Centre, NIH Division of Research (Resources Grant No RR01811). The computations were carried out with the help of Professor F.G. Herring and Dr. P.S. Phillips on an IBM compatible 80386-based personal computer equipped with an Intel 80387 floating point coprocessor.  M6ssbauer Spectrometer The MOssbauer spectra of the dimethyltin(IV) derivatives were recorded at 77 K with the samples placed in Teflon cells on a constant acceleration type spectrometer operating in the transmission mode. The pulses were accumulated on a Tracor-Northern TN-1706 Multichannel Analyzer, and the data were then transmitted via an IBM PC to an IBM 3070 for curve fitting using a non-linear regression program with no restrictions placed on the fitting  49  parameters. The Doppler velocity scale was calibrated using a 57 Co source with an iron foil as an absorber. The isomer shift was calculated relative to BaSnO 3 (and was equivalent to a SnO 2 reference). The error limits for the isomer shifts (S) and the quadrupole coupling (AE Q) are ± 0.03 mm.s -1 . The spectra were obtained with the assistance of Dr. J.R. Sams and Mrs. A. Sallos.  X-Ray Diffractometer The single crystal X-ray diffraction studies were carried out on a Rigaku AFC6S diffractometer using Mo K a radiation of 0.71069A wavelength. Crystals were mounted and sealed off in 0.3 or 0.5 mm i.d. quartz capillary tubes (Charles Supper Company, Ltd., MA) under a nitrogen atmosphere. Further details are available in the relevent chapters following. The crystal structures were obtained with the courtesy of Dr. S.J. Rettig and Dr. J. Trotter of the X-Ray Crystallography Laboratory, Department of Chemistry, UBC.  Mass Spectrometer The mass spectra were obtained on a Kratos MS50/DS55SM (1974) spectrometer using a direct insertion probe that was loaded in a glove bag filled with dry nitrogen. Fragmentation was carried out using electron impact methods, and system software was used for the data analysis. The mass spectra were obtained courtesy of M. Lapawa and Dr. G. Eigendorf, Mass Spectrometry Laboratory, Department of Chemistry, UBC.  Melting_point Determinations The samples were loaded into 2.0 mm o.d. capillary tubes in the drybox and flame sealed off under nitrogen. The melting points were obtained on either a Gallenkamp or Thomas Hoover Unimelt capillary melting point apparatus.  50  MICROANALYSES Elemental microanalyses for carbon, hydrogen, nitrogen, chlorine and sulfur were carried out by Mr. P. Borda of this department. Uranium was analyzed gravimetrically by the pyrolysis of the compound in air at ca. 800°C to quantitatively yield U 3 0 8 . Titanium, zirconium, hafnium, molybdenum, cesium and fluorine were analyzed by either Mikroanalytisches Laboratorium Beller (Gottingen, Germany) or Analytische Laboratorien (Gummersbach, Germany).  GENERAL REAGENTS The reagents described here are the general chemicals used in the course of this study. Those obtained from commercial sources are listed below with some notes on their purity, and the methods used, if any, to purify them further before use. The remainder were synthesized along previously reported precedents. In addition to these compounds, bis(fluorosulfuryl) peroxide (S 2 06 F2 ) was used extensively for the synthesis of all the fluorosulfate derivatives, and its large-scale synthesis is described in some detail below. The remainder of the chemical compounds used throughout the course of this study will be described in the appropriate sections.  Cesium chloride (99.9%), phosphorus pentoxide (98%), calcium hydride (reagent grade), obtained from BDH Chemicals, and sulfur trioxide (Allied Chemicals) were used without further purification. Technical grade fluorosulfuric acid (Orange County Chemicals) was purified by double distillation at atmospheric pressure against a counterflow of dry nitrogen as described before [5], with the constant boiling fraction (162°C) collected in a dry Pyrex storage bulb.  51  SYNTHESIS OF BIS(FLUOROSULFURYL) PEROXIDE (S 2 06 F 2) AND FLUORINE FLUOROSULFATE (FOSO 2F) Bis(fluorosulfuryl) peroxide was synthesized by the catalytic (AgF 2) reaction of F2 on SO 3 in a nitrogen flow system. The original synthesis, [6][7] was adapted, as discussed below, to yield large quantities of S 20 6F 2 in a fairly straightforward manner. The same reaction may be used for the synthesis of fluorine fluorosulfate, FOSO 2 F [8], which requires higher temperatures and a 1:1 F 2 :S0 3 ratio in contrast to S 2 06 F2 which requires a 1:2 ratio. Fluorine fluorosulfate is known to be an extremely hazardous and unpredictable compound [9][10], and is often formed in small quantities as a byproduct in the synthesis of S 2 06 F2 , although lower temperatures and a lower F 2 :S0 3 ratio are used.  The Flow System  The reaction was carried out in a flow system of the type illustrated in Fig. 2-7. The fluorine and the sulfur trioxide were carried by nitrogen streams into the flow reactor, which was held at 180°C. The volume of the lines was calibrated, enabling the quantities of the reactants to be calculated, and the rate of flow was used to control the amounts of the reactants consumed. A heat lamp or a heating tape was used to heat the SO 3 to 40°C to increase the SO 3 throughput and facilitate its transport by the stream of nitrogen. A series of four glass traps with brass adapters was used to trap the reaction products (Fig. 2-7, A-D). Trap A was kept at room temperature to cool the reaction effluents slowly and to allow for visual detection of any unreacted SO 3 that might condense at this temperature. Traps B and C were held at -78°C (powdered dry-ice) to collect the S20 6 F2; the two traps efficiently condense the S 20 6 F2 while allowing the more volatile by-products to escape. Trap D was also kept at -78°C instead of at -176°C as suggested in the original synthesis [7]. This trap was used to collect any residual S 2 0 6 F 2 that was carried over in the N2 stream while allowing any unreacted F2 and the byproduct, FOSO 2 F, to pass through to the soda-lime trap where they were converted to water and a solid residue of CaF2 and Ca(SO3F)2.  52  The Flow Reactor and Catalyst  The tower reactor (Fig. 2-8) consisted of a Monel tube with a flange and a removable top to allow the addition and removal of the catalyst. The original Teflon gasket was destroyed during one preparation when the temperature inside the reactor reached ca. 400°C at the reaction front, possible due to an anaerobic fluorine fire within the reactor. The high temperatures cause Teflon to flow, and distortion of the gasket resulted in a leak at the flange. The Teflon gasket was then replaced by a copper gasket which prevented further unexpected leaks. The Monel tube was wrapped in two Chromel heating coils which were used to control the temperature of the top and bottom half of the reactor. Thermocouple wires were used to measure the temperature which was maintained at 135-150°C. This assembly was surrounded by Fiberfax insulating paper and Pyrex wool, and then encased in 0.02" thick stainless steel plating held down by hose clamps. Two additional thermocouple wires in Monel probes were inserted into the top and bottom of the reactor to measure the internal temperatures and thus follow the reaction more closely. Copper turnings (Johnson and Mathey) were silver plated using a AgNO3 and KCN bath with [Ag(CN) 2r as the principal complex. The turnings were washed and dried after which they were loaded into the reactor and fluorinated in a slow undiluted stream of F2. The temperature was controlled by the flow rate of fluorine gas. After completion of the reaction, the reactor was flushed by a stream of diluted fluorine followed by a dry nitrogen purge during which the reactor was allowed to cool to room temperature. The reactor was then sealed off under a nitrogen atmosphere till further use. The following problems have been encountered at some point during our operations. High internal temperatures greater than 1000°C have caused melting of the copper wool and resulted in a blockage of the flow reactor; along with the blockage, a hole was burnt into the side of the reactor causing leakage of the sulfur trioxide, fluorine and the product. Both problems emphasize the need for closely monitoring the internal temperature and the flow rate of the fluorine.  54  Fig. 2-8. Flow reactor used for the synthesis of bis(fluorosulfuryl) peroxide, 8 2 06 F2 .  55  Product Isolation and Purification The product collected in traps B and C (Fig. 2-7) was further treated to remove two principal impurities — SO 3 and FOSO 2 F. The product was purified by trap-to-trap distillation through three traps held at -10°, -78°C and -150°C, respectively. Portions of the product mixed with SO3 were washed with concentrated H2SO4 to dissolve the SO 3 via liquid-liquid extraction. The purity was checked by infrared, Raman and  19 F  NMR spectroscopy, and the product was  stored in sealed glass ampoules at room temperature or -196°C. Traces of S205F2 were found in the batches treated with H 2 SO4 . This method produced about 2-3 kg of S 2 06 F2 over a 36 hour period. The low-boiling fraction was found to contain considerable quantities of fluorine fluorosulfate. The FOSO 2F was further purified by trap-to-trap distillation, and its purity was confirmed by infrared and 19 F NMR spectroscopy. Small portions of the purified product were stored in glass ampules at -196°C in a long-term storage dewar for use in the reactions described in Chapter 6. Caution: Fluorine fluorosulfate is an extremely unpredictable and hazardous byproduct. Any  quantities of FOSO 2F formed in future preparations of S2O 6 F 2 should not be trapped and stored; the FOSO 2F may be safely disposed of by neutralization with dilute solutions of sodium hydroxide or sodium bicarbonate in a well-ventilated fume-hood with suitable safety precautions.  56  REFERENCES  1. W. Gombler and H. Willner, J. Phys. E. Sci. lnstrum., 20 (1987) 1286. 2. F.G. Herring and P.S. Phillips, J. Magn. Reson., 57, (1984) 43. 3. F.G. Herring and P.S. Phillips, J. Magn. Reson., 62, (1985) 19. 4. R.E.D. McLung, Can. J. Phys., 46 (1968) 2271. 5. J. Barr, R.J. Gillespie and R.C. Thompson, Inorg. Chem., 3 (1964) 1149. 6. F.B. Dudley and G.H. Cady, J. Am. Chem. Soc., 79 (1957) 513. 7. G.H. Cady and J.M. Shreeve, Inorg. Synth., 7 (1963) 124. 8. F.B. Dudley, G.H. Cady and D.F. Eggers, Jr., J. Am. Chem. Soc., 78 (1956) 290. 9. G.H. Cady, Inorg. Synth., 11 (1968) 155. 10. G.H. Cady, Chem. Eng. News, (1966) 40.  57  CHAPTER 3. DIMETHYLTIN(IV) CARBOXYLATES AND ANION BASICITIES  INTRODUCTION As discussed in Chapter 1, ionic self-dissociation equilibria are important in defining the useful range of BrOnsted acids and conjugate superacid systems. The following two general equilibria may be used for:  a) Br6nsted acids: 2 HA  Kid H2A+ + A -  b) Conjugate superacids: Kid  ,  2 HA + Y ^ H2A+ + [AY]"  -2-  The extent of ionic dissociation may be explored by conductometry entailing specific conductivity vs. concentration plots, conductometric titrations (the reverse of Equation 1), and equivalence or molar conductivity and transport number determination. These aspects have been discussed previously [1]. An important quantity is the acceptor ability of the Lewis acid which will affect Kid', and this has also been discussed before [2]. The usefulness of both systems (a and b, above) is best thought of in terms of the levelling concept of non-aqueous solvents which states that "The self-ionization ions, the acidium ion (H 2 A + ) and the base ions (A - or [AY]"), are the strongest acids and bases in the given system, respectively." A highly electrophilic cation X + or HX + can only be generated when the acid strength of H 2 A + is sufficiently high to protonate X to HX + , or form X + by dissociation from a suitable precursor. Both X + or HX + will only be stabilized when the  basicity of A - or [AY]" is sufficiently low to prevent an unwanted destructive side-reaction of the initial type:  58  X + + ^ > XA^  -3-  or an analogous reaction for any of the other species. These reactions are viewed as Lewis acid-Lewis base reactions. There are two general approaches used to determine the acidity of the two systems (BrOnsted acids and conjugate superacids): a) Hammett acidity (-H0) determinations of protonic and conjugate systems [3]; or b) determination of the Lewis acid strength of Y. While the nucleophilicity of A' or [AY]' influences both equilibria (1) and (2) and affects the acidity of BrOnsted or conjugate systems, it has not been possible to measure this quantity for extremely weak nucleophiles such as F, SO 3 F, SbF6 - and [SbF 5 S0 3 F]' separately or even produce a relative scale or a ranking. Such an attempt was undertaken in this research group a few years ago [4]. This approach makes use of the fact that many of the anions form dimethyltin(IV) salts of the type (CH 3 ) 2 SnA 2 or (CH3) 2 Sn[AY] 2 .  Organotin compounds  Organotin compounds contain at least one direct Sn-C bond which is usually covalent. In these compounds, tin predominantly exhibits a valency of four, much like its carbon analogues, although some organotin(II) compounds are known to exist. The annual production of organotins in the mid 1980s was believed to be in excess of 3.0 x 10 7 kg [5]. The relative ease of synthesis of organotin compounds has been instrumental in the amount of research in this field [6]. Organotins are of interest as stabilizers in the polymer industry [7], fungicides in agriculture [8], anti-tumour agents in medicine [9], and catalysts in the vulcanization of silicone rubbers [10]. They are also extensively used in anti-fouling paints for marine craft [11].  59  The coordination sphere around tin can be expanded beyond 4 (usually 5 or 6), unlike that for carbon. This covalency expansion is possibly due to the presence of vacant d-orbitals on tin [12] and is often noted for elements where the principal quantum number n ^3. Alternatively, the Sn(IV) center in compounds of the type R n SnX (4 _„) becomes a good acceptor with decreasing n, and will undergo oligomerization or polymerization via X-bridges under three conditions: (a) in order to minimize steric hindrance, R must not be bulky; (b) X must be electron withdrawing as well as potentially polydentate; and (c) good Lewis bases or donor solvents must be absent. As a general rule, R 4 Sn are tetracoordinate, R3SnX are frequently pentacoordinate and R 2 SnX 2 or RSnX 3 are hexacoordinate. Previous work in our group has been concerned with dimethyltin(IV) derivatives of strong protonic oxyacids or fluoroacids. It is generally accepted now that these salts are not ionic but better viewed as polymers where A" or [AY]" acts as a bridging bidentate ligand. The oxygen or fluorine atoms on the oxy- or fluoro-anions coordinate to tin to give an octahedral coordination about tin, with the C-Sn-C moiety in a trans arrangement. This model, confirmed by a number of studies, can be shown in its two limiting cases as below: z^  z  1^ Y C^2+  t C  (F)0 .,^. 0(F)^ (F)0^ 4-------1 .^ .Sn I (F)0 (F)0 C^  -...A.  x  Limiting case A^  C Limiting case B  When the nucleophilicity of A- or [AY]- approaches zero, we would expect the limiting case of the dimethyltin(IV) cation, (CH 3 ) 2 Sn 2+ . The bonding in this cation involves a spz hybrid on tin and suitable acceptor orbitals on the 0 or F atoms in the groups A' or [AY]"; both the px and py orbitals on tin are empty.  60  Yeats [13] had established simple straightforward synthetic routes and, by means of vibrational spectroscopy and MOssbauer spectroscopy, had reached some conclusions regarding the structures of these derivatives. The synthesis of choice is the acidolysis of derivatives such as (CH3)4Sn, (CH3) 3 SnC1 and (CH3) 2 SnC1 2 , at room temperature, which results in disubstitution on tin, and the formation of CH 4 and HC1. The Sn-Cl bond is cleaved preferentially over the Sn-C bond [14]. The dimethyltin moiety remains intact under these conditions. The acids used so far have been a large variety of strong monoprotonic acids such as HF, HSO 3 R (R = F, Cl, CF 3 and CH 3 ) and HPO 2 R2 (R = F, C 6 H5 ) [14][15][16].  25°C (CH 3 )„SnC1 (4 ,) ^ > (CH 3 ) 2 SnX 2 + (n-2) CH 4 + (4-n) HC1^-4excess HX where X = SO 3 F, SO 3 CF 3 , SO 3 CH 3 , PO 2F 2 , F, etc. and n = 2, 3 or 4.  In general, this method is applicable when the acid in question is a strong acid, liquid at room temperature, and acts as both an ionizing solvent and a reagent. The alkyl chain is best kept short so as to avoid unwanted side reactions resulting in discoloured products. Conjugate superacid media may also be used as reagents. The compounds obtained may be viewed simplistically as salts where the linear organometallic cation (CH 3 ) 2 Sn 2+ is stabilized by acid (or superacid) anions, in which the covalent cation-anion interactions increase with increasing basicity of the counteranions. The structures of (CH 3 ) 2 SnF 2 [17] and (CH 3 ) 2 Sn(SO 3 F) 2 [18] show some interesting similarities. Both compounds are sheet polymers with the central tin atom coordinated to either four fluorine or four oxygen atoms arranged equatorially in a plane. The linear CH 3 -Sn-CH3 group is perpendicular to the plane formed by the symmetrically bridging F - or SO 3 F - anions. The Sn-F and Sn-O distances are quite long, and the SO 3 F bond distances and  61  angles, except for a slightly shorter S-F bond, are comparable to those found in KSO3F [19]. These features suggest an ionic structure in which the (CH3)2Sn 2+ cations are packed perpendicular to a sheet formed by bridging bidentate SO3F - ions (Fig. 3-1).  Bond Distances (A) Sn - C  , • 0 Sn S^0  F C  Fig. 3-1. Crystal structure of (CH 3 ) 2 Sn(SO3F)2 [18]  Sn-O(1) Sn-O(2) S-0(1) S-0(2) S-0(3) S-F  2.065(9) 2.270(6) 2.271(7) 1.437(8) 1.428(7) 1.422(11) 1.504(8)  62  This type of symmetric structure is not seen in all dimethyltin(IV) derivatives. The structure of (CH 3 ) 2 SnC1 2 [20][21] indicates that the C-Sn-C angle is 123.5°, and the chlorine bridges are asymmetric, resulting in a distorted environment about the central tin atom. The [(CH 3 ) 2 SnC1 4 ] 2- anion, however, exhibits a linear (CH 3 )2Sn grouping and a symmetric Sn-Cl coordination [21]. Other derivatives such as (CH 3 ) 2Sn(CN)2 [22], (CH 3 ) 2 Sn(NCS) 2 [23] and (CH 3 ) 2 Sn(Mo04) [24] also show the presence of asymmetric bridging anions and a significant departure from linearity for the C-Sn-C grouping.  M6ssbauer Spectroscopy The nuclear transformation of  119mSn  (119 metastable tin) to 119 Sn proceeds in two steps with  two different quanta of 7-rays emitted. The second quantum may be reabsorbed by naturally occurring 119 Sn (8.58% natural abundance), which exhibits the MOssbauer effect and allows us to obtain well-resolved spectra without having to enrich the samples. The MOssbauer effect and its application in chemistry have been reviewed by Greenwood [25]. This section will review the basic principles of MOssbauer spectroscopy, and discuss its application in structural and bonding problems in dimethyltin(IV) derivatives of the type studied in this work.  Basic Principles The recoilless emission and resonant re-absorption of 7-radiation gives rise to the MOssbauer effect (also known as the 7-ray resonance effect). The process involves transitions between the nuclear energy levels of the emitter and the absorber of the radiation. ground state via two successive 7-emissions (Fig. 3-2a).  119m Sn  119m Sn  decays to its  has a half-life of 245 days.  In the first decay, the nucleus relaxes to the first excited nuclear state (1=3/2). This is followed by a rapid second decay as the nucleus relaxes to the ground state (I=1/2). The transition of interest is that between the ground state and the first excited state of  119 Sn.  The half life of the 1=3/2 first excited state (1.84 x 10 -8 s) is in a range where the linewidths are at an optimum. The source is a polymeric molecular substance (Ba l 19mSn03)  63  a)^ 119  mSn  ^t 12 = 245 d  y, = 65.66 keV  t ia = 1.84 x 10 -8 s + 3/2 y,, = 23.875 keV + 1/2 "9  Sn  b)  to memory t  waveform generator  detector 4' to oscilloscope electromechanical transducer and velocity sensor absorber in Dewar  Fig. 3-2. a) Decay scheme for  119m  Sn^b) Block Diagram of MOssbauer Spectrometer  64  embedded in a 7-inert heavy-metal matrix that allows recoilless emission. The parent isotope has a half-life of 245 days, having the advantage of being a long lived source (several years) depending on its original activity. The energy of the second -y-decay (7 m ) is of a magnitude (23.875 keV) where we can study the MOssbauer effect at temperatures as high as liquid nitrogen temperature (77 K) and be assured of recoilless absorption. The block diagram of the MOssbauer spectrometer is shown in Fig. 3-2b.  Chemical Application of the MOssbauer Effect The spacings of the nuclear energy levels are critically dependent on the electron distribution in the 5s and 5p valence shells in tin. This electron distribution is, in turn, influenced by the chemical environment around the nucleus. When the source and the absorber are different, their 7-energies will be different and there will be no absorption. By moving the 7-ray source relative to the absorber, the frequency of the 7-radiation may be modulated by exploiting the Doppler effect, and the energies of the source and the absorber may then be brought into resonance. The M6ssbauer spectrum obtained is a plot of the 7-ray counts against the relative velocity of the source with respect to the absorber (sample of interest), the peaks occurring in the regions where resonance absorption occurs and the 7-ray count drops. The two principal MOssbauer parameters, the isomer shift (b) and the quadrupole coupling (AE Q ), are both measured in terms of the Doppler velocity (mm•s .1 ), and are of interest in structural analysis.  Isomer shift: The isomer shift, S (mm.s 1 ), is the relative velocity of the source, with respect -  to the sample, that is required to bring the energies of the two in resonance. It may be simplistically explained by:  = k•(Ar)/r• th(0) 1 2 where Ar = rexcited rground = difference in nuclear radii I lk s (0) I 2 = difference in s electron density between the nucleus of the source and sample (see Fig. 3-3).  65  Isomer Shift, 8 (mm•s -1 ) MOssbauer spectrum  Absorption %  -4--IS -> 8 1 -2 -1^0^+1^2^3^4 Velocity in mm•s"'  Excited State  Ea  Eo  ■I'  Ground State Source^  Absorber  IS = E. - Es _ 4_g_ 2 2 _EL 5 Ze r •• r lif(0)1 2 absorber - 11140)1 2 source ] -  = const. • -6-r- • 8 1 Ws (0)1 2  Fig. 3-3. Diagrammatic scheme showing the isomer shift  66  The factor (Or)/r has a positive value for tin, and thus a large value for 6 implies a high s electron density around the tin nucleus. One may therefore get an idea of the oxidation state of tin in the sample. The isomer shift scale is arbitrarily set up so that the value of the isomer shift is 0 mm.s -1 for the reference (SnO 2 (5s0)), 2.1 mm•s -1 for grey tin (5s 1p 3 ) and 3.7 mm.s -1 for SnC1 2 (5s2 ). With the exception of binary Sn(IV) fluoro and fluorosulfato derivatives, all tin compounds have positive isomer shift values. The isomer shift (6) relative to Sn02, thus becomes a measure of the s electron density on tin and would be the highest for (CH 3 )2Sn 2± .  Quadrupole Splitting: The first excited state of 119 Sn (I = 3/2) posseses a nuclear quadrupole moment whose orientation depends on the electric field gradient, q. If the tin nucleus has an asymmetrical electron distribution around it, the electric field gradient thus produced in the molecule removes the degeneracy of the I=3/2 state resulting in quadrupole splitting LE Q . This splits the peaks (see Fig. 3-4) in the spectrum as explained by: LEQ = (1/2)e 2 •q•Q•[1 + (1/3)n 2] I /2 = k•q•Q•(1 + (1/3)17 2 ] 1/2 where,^k = constant, q = electric field gradient,  Q = nuclear quadrupole moment, n = asymmetry factor. In dialkyltin(IV) compounds which exhibit axial symmetry,  n = 0, and we get:  AEQ = k•q•Q  The electric field gradient in tin(IV) compounds arises mainly from the polarity of the a-bonds, and the major contribution to the electric field gradient is the imbalance of the valence shell electron distribution around tin, [q vd]. The magnitude and sign of [q va d reflect the orientation of the ligands around the tin atom and the imbalance in the polarity of the tinligand bonds. The magnitude of the quadrupole splitting is proportional to the difference in  67  Quadrupole Coupling, A (rnm.s -1 )  Absorption %  r^I  I^I^I  0  8  Velocity in mm•s 4  I Excited state  3 I —^2  Ground state  1  I= 2  m ^ + a t^- 2 ^ + 1 -  2  +1 Isomer^Quadrupole - 2 shift^coupling  Fig. 3-4. Diagrammatic scheme for quadrupole splitting.  68  the bond polarities of the ligands, for compounds having the same geometry around the tin nucleus. Thus the quadrupole splitting (AE Q) expresses an imbalance in the p electron distribution (p2 vs. px and py) in the tin orbitals. In the limiting case B (vide supra, page 59) where tin forms six covalent bonds to C and 0 (or F), the bonding could be described as  sp 3 d2 . Assuming the retention of a linear C-SnC grouping and a symmetrical coordination environment about tim, increasing the nucleophilicity of the anions will increase the electron population of the 5px and the 5py orbitals, both of which are directly oriented towards the F (or 0) atoms of the ligands. This will reduce the quadrupole splitting as well as the isomer shift (due to enhanced screening). As one goes from the limiting case A towards the limiting case B, the isomer shift decreases gradually along with the quadrupole splitting; both would reach their lowest values at case B. Hence, we would expect high values for 6 and AE Q for those (CH 3 ) 2 SnA 2 compounds in which the anions are weak nucleophiles. Indeed, a quadrupole splitting of 5.55 mm.s -1 was found for (CH 3 ) 2 Sn(SO 3 F) 2 [16], and (CH 3 ) 2 Sn[SbF 6] 2 and (CH 3 ) 2 Sn[Sb 2F i1 ]2 have quadrupole splittings slightly greater than 6.00 mm.s -1 [26]. These fluoroantimonate derivatives have exhibited isomer shifts slightly greater than 2.00 mm•s -1 , the highest isomer shift found for any organotin(IV) compounds [26]. The advantages of 119 Sn MOssbauer spectroscopy may be summed up as follows: a) The natural abundance (8.58%) of  119 Sn  is sufficient to give a measurable effect even in  samples with a high molecular mass and a low tin content, without having to resort to isotope enrichment. b) The half-life of  119 Sn  (245 days) differs considerably from those of other  MOssbauer-active isotopes, which are often on the order of a few days. This is of practical importance, because  119m Sn  sources of 5-10 mCurie activity have a useful life of about 3-5  years before low activity necessitates longer measurement periods of 24-48 hours rather than 4-6 hours. c) At 23.875 keV, y m is quite low, and the source may be operated at ambient temperatures with a sufficiently high recoilless fraction to allow reabsorption with the solid sample at 77-300 K. d) The natural linewidth of MOssbauer lines (0.626 mm.s -1 ) — a direct  69  consequence of the lifetime of  119m Sn  — is sufficiently small to give sharp lines. e) The  I = 1/2 to I = 3/2 transition produces simple spectra. Single line spectra are caused by tin in a high-symmetry environment (Td as in SnR 4 , or Oh as in SnF 62-), and doublets are caused by a non-zero quadrupole splitting (vide supra).  Anion Basicities As discussed in Chapter 1, the limits of BrOnsted-Lewis conjugate superacids are governed by the basicity of the anion X - or EX (n+ 0 - in:  2 HX + EX n  > H 2 X + + EX (n+i) -^-5-  The anion basicity may be measured by using the dimethyltin(IV) cation as a M6ssbauer probe. For a series of strong protonic fluoroacids and oxyacids, a linear relationship has been observed for the MOssbauer parameters (5 and AE Q , where the isomer shifts range between 1.2 and 2.1 mm-s -1 , and the quadrupole splittings vary from 4.0 to 6.1 mm. s -1 (see data in Table 3.1). From the linear plot (Fig. 3-5) it can be seen that the relationship may be expressed as:  .5 = 0.55(AEQ - 2.21) mm•s -1 or AEQ = 1.82.6 (rel. Sn0 2 ) + 2.21 mm•s -1 The position on the graph, therefore, gives us an empirical measure of anion basicity. As the correlation in the graph shows, this concept works well for anions of strong acids such as SO3R - (R = F, Cl, CF 3 , CH3, etc) or MF 6- (M = Nb, Ta and Sb) and also for PO2R2 (R = F, CH 3 ). As discussed earlier, salts of the weaker acids such as (CH3)2SnC1 2 [20], (CH 3 )2Sn(CN) 2 [22], (CH 3 ) 2Sn(NCS) 2 [23] and (CH3)2SnMoO4 [24] show distorted geometries about the Sn centre and, hence, cannot be incorporated into the linear correlation above. While the coordinative asymmetry of bridging groups such as CN or  70 NCS may be viewed as a plausible cause for the observed distortions, this explanation is, however, not valid for the MoO 4 group. Table 3.1.  119 Sn MOssbauer data of some dimethyltin(IV) derivatives of strong protonic acids  and superacids.  Compound  (5 onnvs -ir,b  AEQ (mm.s -1 ) a  Ref.  (CH 3 ) 2 SnF 2  1.23  4.52  27  Li 2 [(CH 3 ) 2 SnF 4 ]  1.31  4.60  28  (CH 3 ) 2 Sn(S0 3 -p-CH3C 6 H 4 ) 2  1.51  4.85  16  (CH 3 ) 2 Sn(SO 3 C 2 H5) 2  1.55  4.91  16  (CH 3 ) 2 Sn(SO 3 CH3)2  1.52  5.05  16  (CH 3 )2Sn(P02F2)2  1.53  5.13  14  (CH3)2Sn(S030)2  1.75  5.20  16  (CH3)2Sn [TaF6]2  1.69  5.23  26  (CH3)2Sn[NbF6]2  1.77  5.51  26  (CH3)2Sn(SO3CF3)2  1.79  5.51  16  (CH 3 ) 2 Sn(SO 3 F) 2  1.80  5.55  16  (CH3)2Sn[Pt(S03F)6]  1.95  5.67  29  RCH3)2SnS0392[Pt(SO3F)6]  1.96  5.70  29  (CH 3 )2Sn[Sb2F1 1 1 2  2.08  6.01  26  (CH 3 ) 2 Sn[SbF 6 1 2  2.04  6.04  26  'confidence limit ± 0.03 mm•s -1 ; brelative to SnO2 at 77 K. The molecular structures of three dimethyltin(IV) carboxylates of the type (CH 3 ) 2 Sn(O 2 CR)2 (R = H, CH 3 and CF 3 ) were determined during the course of the present  71 6.50  6.00 -  5.50 _-  AE (  ^  .s-i ) 5.00 =  4.50 =  4.00 1.00^1.20^1.40^1.60^2.00^1.80^2.20  5 (mm•s I )  Fig. 3-5. Correlation of isomer shift and quadrupole splitting for dimethyltin(IV) salts of strong fluoro- and oxyacids.  study to gain a better understanding of why substantially distorted geometries are encountered when anions of the weaker acids are involved. The above compounds were selected for this study in order to minimize steric repulsion induced by the R group on the RCO2. An extension of the above correlation to the weaker protonic acids of the type RCO 2 - is also attempted in this thesis. Three interrelated requirements must be met by the dimethyltin(IV) derivatives for the model to be useful: a) A - or [AY] - must be capable of functioning as a bidentate ligand. A preference for bridging rather than chelating coordination is desirable.  72  b) A or [AY] must function as symmetrically bridging ligands. This implies that all the -  -  Sn-O(F) distances must be identical within estimated standard deviations. c) The C-Sn-C group must remain linear.  EXPERIMENTAL Reagents Dimethyltin(IV) dichloride (reagent grade) was obtained from either Aldrich Chemical Co. Inc. or K & K Laboratories, Inc., and purified by sublimation in vacuo. The following reagents were used without further purification: Dimethyltin(IV) oxide (reagent grade), from Alfa Inorganics, acetic acid (99.7% purity) and acetic anhydride (97% purity) from Amachem, formic acid (98% purity) and trifluoroacetic acid (99.5% purity) from BDH Chemicals, and trifluoroacetic anhydride (Gold Label, 99%+ purity) from Aldrich Chemical Co. Inc.  Reactions While some the compounds in this work have been prepared previously, the experimental details are included here to explain to the reader the methods used for obtaining single crystals suitable for a diffraction study.  Dimethyltin(IV) bis(formate) was synthesized as reported previously [30]. About 270 mg of (CH 3 ) 2 SnO were dissolved in an excess (ca. 5 mL) of formic acid. The mixture was refluxed for 4 days, and then allowed to cool to room temperature. After 7 days, clear crystals were observed at the bottom of the reactor.  Dimethyltin(IV) bis(acetate) was synthesized as reported previously [31]. About 200 mg of (CH 3 ) 2 SnO were dissolved in ca. 5 mL of a 1:1 mixture of (CH 3 C0) 2 0 and CH3CO 2 H. The reaction mixture got warm and a clear, pale-yellow solution was obtained. This was allowed  73  to stir under vacuum for 1 d and then left sealed under nitrogen. After a period of about 4 months, colorless prismatic crystals were observed in the reaction vessel.  Dimethyltin(IV) bis(trifluoroacetate) was synthesized as using an adaptation of the method reported previously by Okawara et al. [30][31]. About 200 mg of (CH 3 ) 2 SnO were dissolved in ca. 5 mL of a 1:1 mixture of (CF 3 C0) 2 0 and CF3CO 2 H. The reaction mixture initially got warm to the touch, and was left stirring at room temperature for 12 h. The clear pale-yellow solution was shown signs of turbidity. The reaction mixture was immersed in a hot water bath at 65°C and allowed to cool slowly to room temperature while in the water bath. After three days, a few irregular crystals were formed in the reaction vessel.  DISCUSSION Structural information on many dimethyltin(IV) salts has been reported over the years; however, there is little information on the carboxylate derivatives. Although the diorganotin(IV) carboxylates are widely used (vide supra), the molecular structures of these compounds are open to speculation. While some of these salts were first synthesized in the late 1950s [30], there are several unanswered questions about the coordination number of tin, the molecularity of these compounds and the coordination mode of the carboxylate anion. The extreme hydrolytic instability of these compounds has reportedly hindered structural investigations in some studies [30][32]. There are several conflicting structural proposals in the literature based on investigations involving infrared [31][33], NMR [34], and MOssbauer [35] spectroscopy. The are several possible structures for the diorganotin(IV) carboxylates. The presence of only chelating carboxylates results in monomers, while the presence of both chelating and bridging ligands can give rise to dimers or oligomers. The presence of bridging groups alone can result in either chains, sheets or three-dimensional network polymers. Various spectroscopic studies (vide supra) on dimethyltin(IV) carboxylates are often in conflict and  74  indicate that the solution structures of many of these compounds are either monomeric or dimeric. NMR investigations by Lockhart et al. [32] indicate that the solution and solid-state structure of (CH 3 ) 2 Sn(O 2CCH 3 ) 2 is the same, and the C-Sn-C angle of the dimethyltin moiety is calculated to be ca. 135°. Based on available data, they strongly favoured the monomeric structure as being representative of (CH 3 ) 2 Sn(O 2 CCH3 )2 and other diorganotin(IV) carboxylates [36]; this proposal is partially supported by the X-ray crystal structure of (CH3) 2 Sn(O2CCH 3 )2 [36]; this crystal structure, obtained twenty years after the synthesis of the compound, was the first X-ray crystal structure of a diorganotin(IV) carboxylate ever to be reported. The details of the crystal structures obtained in the present study are listed in Table 3.2 below. The (CH 3 ) 2 Sn(O2CCH 3 ) 2 prepared in the present study has a structure similar to that reported by Lockhart et al. [36], but is observed to exist in a different crystal form as is discussed later. The structure of (CH 3 ) 2 Sn(O 2 CH)2 (See Fig. 3-6) consists of infinite sheets with symmetrically bridging formate groups linking linear (CH 3 ) 2 Sn moieties. The coordination environment about the Sn atom is that of a relatively undistorted octahedron. There is exact mirror symmetry with the C(1), Sn and C(2) of the (CH 3 ) 2 Sn moiety lying in the mirror plane, and Sn-O = 2.246(2) and 2.249(1)  A, Sn-C = 2.097(3) and 2.116(2) A and  C-Sn-C = 179.7(2)°. The structure falls into the same category as those of (CH 3 ) 2 SnF2 [17] and (CH 3 ) 2 Sn(SO 3 F)2 [18] (see Introduction) although the Sn-C distances in the (CH3 ) 2 Sn(O 2 CH) 2 (ay. 2.107 A) are longer than those reported for (CH 3 ) 2 SnF 2 and (CH 3 ) 2 Sn(SO 3 F) 2 (2.08(1) A). The difference in the distances could be attributed a combination of: a) the intralayer steric interactions (2.31 A) between the H(5) hydrogen bonded to the formate and the H(1) hydrogen of the methyl group; and b) the electronic effect of the more acidic SO 3 F group in the fluorosulfate derivative. The Sn-O bonds in the formate (ay. 2.248 A) are shorter than those in the fluorosulfate (ay. 2.270 A), which also reflects the higher basicity of the formate anion as compared the the fluorosulfate.  75  Table 3.2. Crystallographic data for (CH 3 ) 2Sn(O 2 CR)2 derivatives  (CH3 ) 2Sn(OOCH)2  (CH3)2Sn(OO C CH3)2  (CH 3) 2Sn(OOCCF3) 2  238.79  266.85  374.79  Sheet  Monomer  3D-polymer  orthorhombic  monoclinic  monoclinic  Pnma (#62)  C2/c (#15)  Cc (#9)  a (A)  12.693(2)  26.282(4)  8.444(1)  b (A)  9.128(2)  5.282(1)  17.689(1)  c (A)  6.013(2)  14.434(3)  15.368(1)  101.17(2)  90.013(9)  696.7(2)  1966(1)  2292.3(4)  4  8  8  2.28  1.803  2.172  0.030; 0.037  0.025; 0.023  0.027; 0.030  179.7(2)  133.8(2)  149.3(2)  formula wt. Nature Crystal system Space group  ft (°) V (A 3 ) Z  Pcalc (g.m1--1) .  R; R H,  C-Sn-C (°)  The Sn-O bonds involving the 1.1.-0 2 CH are different even though the two Sn-O distances are essentially equal. The Sn-O(1) bond is cis to the H(5) atom on the formate while the Sn ii-0(2) is trans to the H(5). The Sn and Sn # atoms are out of the plane of the formate ligand by -0.337 and -0.255  A. The H(5)-H(5) separation of 2.58 A in this structure suggests  that other compounds of the type (CH3)2Sn(O2CR)2 where R H would not have the same structure as that described for the formate. The bond lengths and angles for [(CH 3 ) 2 Sn(02CH)2L are given below (see Table 3.3) along with their estimated standard deviations.  ^-  76  Sn#  Sn#^  11'  •^ 40•  •  .^it •^•  14 •^lb  •  • IP  aL •  ,  11•1111111111■1 , ■ 4  ^dip  • •  •  • lib •^•^•  Fig. 3-6. a) Stereoview down the  a axis of the structure of [(CH 3 ) 2 Sn(02 CH) 2L  b) Packing diagram of [(CH3) 2 Sn(02CH)2]„ viewed along the a axis  77  Table 3.3. Bond lengths (A) and angles (°) for (CH 3 ) 2Sn(O2 CH) 2 with estimated standard deviations in parentheses  Sn-C(1)^2.116(2)  Sn-0(1)  2.246(2)  0(1)-C(3)^1.249(2)  Sn-C(2)^2.097(3)  Sn-0(2)'  2.249(1)  0(2)-C(3)^1.254(2)  C(2)-Sn-C(1)  179.7(2)  0(1)-Sn-0(2)'  92.78(7)  C(2)-Sn-0(1)  86.09(7)  0(1)"-Sn-0(2)'  173.07(5)  C(2)-Sn-0(2)'  87.51(7)  0(2)'-Sn-0(2) *  84.29(8)  C(1)-Sn-0(1)  94.15(7)  C(3)-0(1)-Sn  129.6(1)  C(1)-Sn-0(2)'  92.23(8)  C(3)-0(2)-Sn #  128.5(1)  0(1)-Sn-0(1)"  89.4(1)  0(1)-C(3)-0(2)  125.7(2)  The structure of (CH 3 ) 2 Sn(O 2 CCH 3 ) 2 (see Fig. 3-7) is made of monomers separated by normal van der Waals contacts. The coordination environment about the Sn atom is that of a highly distorted octahedron and, as in the case of the structure reported by Lockhart  et al. [36], may also be decribed as a bicapped tetrahedron with the 0(2) and 0(4) atoms capping the faces. The C-Sn-C angle is 133.8(1)°, compared to 135.9(2)° reported previously [36]. The acetate groups are anisobidentate, (ay. Sn-0 = 2.096 and 2.58 A) and are nearly coplanar with the tin atom, and the methyl groups on the acetate are 2-fold orientationally disordered. At first glance it appears that the structure of (CH 3 ) 2 Sn(O 2 CCH 3 ) 2 obtained in the present study is the same as that obtained previously [36]; however some subtle differences do surface upon closer examination of the data. There appear to be two crystal forms of (CH 3 ) 2 Sn(O 2 CCH3)2, although the monomeric units in the two structures are similar (see Table 3.4). The previously reported structure has a unit cell that is half the size of the one obtained in the present study, and the asymmetric unit consists of half the molecule; the  78  C  Fig. 3-7. a) Stereoview of the structure of (CH3)2Sn(O 2 CCH 3 ) 2 b) Packing diagram of (CH3)2Sn(O2CCH 3 ) 2 viewed along the b axis  79 Table 3.4. Bond lengths (A) and angles (°) for (CH3) 2 Sn(O 2 CCH 3 ) 2 with estimated standard deviations in parentheses *  Sn(1)-0(1)  2.098(2)  Sn(1)-0(3)  2.094(2)  Sn(1)-O(2)  2.563(2)  Sn(1)-O(4)  2.595(2)  Sn(1)-C(1)  2.096(3)  Sn(1)-C(2)  2.088(3)  [2.106(2)] [2.539(2)] [2.098(3)]  0(1)-C(3)  1.291(3)  0(3)-C(5)  1.291(3)  O(2)-C(3)  1.230(3)  0(4)-C(5)  1.216(4)  C(3)-C(4)  1.486(4)  C(5)-C(6)  1.487(4)  [1.293(3)] [1.237(3)] [1.496(4)]  0(1)-Sn(1)-0(2)  54.56(7)  O(4)-Sn(1)-C(2)  88.6(1)  0(1)-Sn(1)-0(3)  80.03(8)  C(1)-Sn(1)-C(2)  133.8(1)  0(1)-Sn(1)-0(4)  133.77(7)  Sn(1)-0(1)-C(3)  103.2(2)  0(1)-Sn(1)-C(1)  106.5(1)  Sn(1)-0(2)-C(3)  82.8(2)  0(1)-Sn(1)-C(2)  108.3(1)  Sn(1)-0(3)-C(5)  104.3(2)  0(2)-Sn(1)-0(3)  134.56(7)  Sn(1)-0(4)-C(5)  82.5(2)  0(2)-Sn(1)-0(4)  171.66(7)  0(1)-C(3)-0(2)  119.4(3)  0(2)-Sn(1)-C(1)  88.8(1)  O(1)-C(3)-C(4)  117.6(3)  0(2)-Sn(1)-C(2)  87.5(1)  0(2)-C(3)-C(4)  123.0(3)  0(3)-Sn(1)-0(4)  53.75(7)  0(3)-C(5)-0(4)  119.5(3)  0(3)-Sn(1)-C(1)  107.6(1)  0(3)-C(5)-C(6)  117.4(3)  0(3)-Sn(1)-C(2)  107.5(1)  0(4)-C(5)-C(6)  123.1(3)  0(4)-Sn(1)-C(1)  88.6(1)  * bond distances in square brackets are quoted from Lockhart et al. [36]  80  molecule itself lies on a 2-fold axis, and hence the bond distances and angles are related by symmetry. In the present work, the whole molecule is the asymmetric unit, and while it does not lie on a 2-fold axis, it is close to a 2-fold axis. The bond distances and angles are comparable (see Table 3.4) to those found by Lockhart et al. [36]. The differences in the bond parameters and angles are noticeable, in particular, in parameters that involve the 2-fold axis passing through the molecule in the previously reported structure. A possible cause for the slight differences in the crystal forms may be the manner in which the single crystals were obtained. In the present study, the single crystals were obtained from a CH 3 CO 2H(CH 3 C0)20 solution, while the single crystals obtained by Lockhart et al. [36] were recrystallized from a dilute solution in hexane. The structure of [(CH 3 ) 2 Sn(0 2 CCF 3 ) 2]„ (see Fig. 3-8) is composed of a threedimensional polymer formed by bridging trifluoroacetate groups around an octahedrally coordinated Sn atom. The bond lengths and angles for [(CH 3 ) 2 Sn(0 2 CCF 3 ) 21,, are given below (see Table 3.5) along with their estimated standard deviations. The C-Sn-C angle is 149.3(2)°, and shows significant departure from linearity. The O 2 CCF 3 groups bridge asymmetrically with two short Sn-O bonds (ay. 2.1 A) that are 85° apart and two long Sn-O bonds (ay. 2.6 A) that are 116° apart. Half of the CF 3 groups are found to be 2-fold orientationally disordered. It is observed (see Fig. 3-8) that the CF 3 groups that are disordered are in a cavity in the crystal lattice, and are free to rotate. It appears that the other CF 3 would generate a short F-F contact upon rotation, and are therefore "locked" in place. The three structures shown here are different to one another in terms of the environment about the Sn atom, the molecular nature of the compound and the bridging ability of the carboxylate anion. They explain, to some extent, the controversy regarding the proposed structures of the dicarboxylates that exists in the literature [30]-[35]. Of these compounds, the [(CH 3 ) 2 Sn(0 2 CH) 23 n is the only compound that adheres to the criteria required to fit the correlation of the MOssbauer parameters. The asymmetrical environment  81  •  C1 •  SP  I  04^Sn1 03  •  e  •  01  C4 Fl F3  •  C10^06  '14 07  oe  C12  Cl 1 F12 F12B  Fl OB F10  Fl 1  Fig. 3-8. a) Stereoview of the structure of [(CH3)2Sn(02CCF 3 ) 21, b) Packing diagram of [(CH3)2Sn(02CCF 3 ) 2]„ viewed down the a axis  82 Table 3.5. Bond lengths (A) and angles (°) for RCH 3 ) 2 Sn(0 2 CCF3 )0„ with estimated standard deviations in parentheses *  Sn(1)-O(1)^2.102(4)  F(2)-C(4)  1.277(8)^F(12B)-C(12)  1.17(2)  Sn(1)-0(3)^2.104(3)  F(3)-C(4)  1.301(8)^O(1)-C(3)  1.270(6)  Sn(1)-O(6) 1^2.700(4)  F(4)-C(6)  1.309(7)^0(2)-C(3)  1.212(6)  Sn(1)-O(8) 2^2.619(4)  F(5)-C(6)  1.288(7)^0(3)-C(5)  1.265(6)  Sn(1)-C(1)^2.091(5)  F(6)-C(6)  1.33(1)^0(4)-C(5)  1.225(6)  Sn(1)-C(2)^2.098(6)  F(7)-C(10)  1.279(7)^0(5)-C(9)  1.273(6)  Sn(2)-0(2)^2.585(4)  F(8)-C(10)  1.32(1)^0(6)-C(9)  1.216(7)  Sn(2)-0(4) 3^2.697(4)  F(9)-C(10)  1.301(7)^0(7)-C(11)  1.256(6)  Sn(2)-O(5)^2.101(4)  F(10)-C(12)  1.27(2)^0(8)-C(11)  1.224(7)  Sn(2)-0(7)^2.122(3)  F(10B)-C(12)  1.42(2)^C(3)-C(4)  1.546(7)  Sn(2)-C(7)^2.082(5)  F(11)-C(12)  1.31(2)^C(5)-C(6)  1.532(7)  Sn(2)-C(8)^2.085(5)  F(11B)-C(12)  1.27(2)^C(9)-C(10)  1.550(8)  F(1)-C(4)^1.33(1)  F(12)-C(12)  1.26(1)^C(11)-C(12)  1.522(8)  0(1)-Sn(1)-0(3)  87.0(2)  0(7)-C(11)-C(12)  110.5(5)  0(1)-Sn(1)-0(6) 1  162.6(2)  0(8)-C(11)-C(12)  120.6(5)  0(1)-Sn(1)-0(8) 2  80.2(1)  F(10)-C(12)-F(11)  110(2)  0(1)-Sn(1)-C(1)  92.3(2)  F(10)-C(12)-F(12)  110(2)  0(1)-Sn(1)-C(2)  110.8(2)  F(10)-C(12)-C(11)  113(1)  0(3)-Sn(1)-0(6) 1  77.6(2)  F(10B)-C(12)-F(11B)  102(1)  0(3)-Sn(1)-0(8) 2  164.6(1)  F(10B)-C(12)-F(12B)  114(2)  0(3)-Sn(1)-C(1)  107.3(2)  F(10B)-C(12)-C(11)  104(1)  83  Table 3.5. Bond lengths (A) and angles (°) for [(CH3) 2 Sn(0 2 CCF 3 )2],, with estimated standard deviations in parentheses * (...cont'd)  94.1(2)  F(11)-C(12)-F(12)  97(2)  O(6) 1 -Sn(1)-O(8) 2  116.1(1)  F(11)-C(12)-C(11)  110(1)  O(6) 1 -Sn(1)-C(1)  84.8(2)  F(11B)-C(12)-F(12B)  107(2)  O(6) 1 -Sn(1)-C(2)  78.7(2)  F(11B)-C(12)-C(11)  113(1)  0(8) 2 -Sn(1)-C(1)  81.8(2)  O(5)-Sn(2)-C(8)  108.9(2)  0(8) 2 -Sn(1)-C(2)  82.6(2)  O(7)-Sn(2)-C(7)  107.1(2)  C(1)-Sn(1)-C(2)  149.3(2)  O(7)-Sn(2)-C(8)  95.7(2)  O(2)-Sn(2)-O(4) 3  116.0(1)  C(7)-Sn(2)-C(8)  146.6(2)  O(2)-Sn(2)-O(5)  82.9(1)  Sn(1)-O(1)-C(3)  129.7(4)  O(2)-Sn(2)-O(7)  166.6(1)  Sn(2)-O(2)-C(3)  152.2(4)  O(2)-Sn(2)-C(7)  79.0(2)  Sn(1)-0(3)-C(5)  127.5(3)  O(2)-Sn(2)-C(8)  84.3(2)  Sn(2)4-O(4)-C(5)  150.5(3)  O(4) 3 -Sn(2)-O(5)  160.0(1)  Sn(2)-O(5)-C(9)  126.4(3)  O(4) 3 -Sn(2)-O(7)  77.1(1)  Sn(1)5-0(6)-C(9)  150.2(4)  O(4) 3 -Sn(2)-C(7)  80.8(2)  Sn(2)-0(7)-C(11)  127.4(3)  O(4) 3 -Sn(2)-C(8)  80.9(2)  Sn(1)6-0(8)-C(11)  152.0(4)  O(5)-Sn(2)-O(7)  84.4(2)  O(1)-C(3)-O(2)  129.1(5)  O(5)-Sn(2)-C(7)  97.6(2)  O(1)-C(3)-C(4)  112.1(5)  F(4)-C(6)-C(5)  113.0(5)  O(2)-C(3)-C(4)  118.8(5)  F(5)-C(6)-F(6)  108.0(7)  F(1)-C(4)-F(2)  104.9(7)  F(5)-C(6)-C(5)  113.6(5)  F(1)-C(4)-F(3)  104.9(6)  F(6)-C(6)-C(5)  109.4(6)  F(1)-C(4)-C(3)  109.4(6)  O(5)-C(9)-O(6)  129.6(5)  F(2)-C(4)-F(3)  111.6(7)  0(3)-Sn(1)-C(2)  84  Table 3.5. Bond lengths (A) and angles (°) for RCH 3 ) 2 Sn(02CCF3)2], with estimated standard deviations in parentheses * (...cont'd)  O(5)-C(9)-C(10)  112.3(4)  F(2)-C(4)-C(3)  113.7(5)  O(6)-C(9)-C(10)  118.1(4)  F(3)-C(4)-C(3)  111.7(5)  F(7)-C(10)-F(8)  104.8(6)  O(3)-C(5)-O(4)  128.6(5)  F(7)-C(10)-F(9)  110.8(6)  O(3)-C(5)-C(6)  111.9(4)  F(7)-C(10)-C(9)  111.4(5)  0(4)-C(5)-C(6)  119.5(5)  F(8)-C(10)-F(9)  106.8(6)  F(4)-C(6)-F(5)  108.3(6)  F(8)-C(10)-C(9)  109.3(5)  F(4)-C(6)-F(6)  104.0(6)  F(9)-C(10)-C(9)  113.3(5)  F(12)-C(12)-C(11)  116(1)  0(7)-C(11)-0(8)  128.8(5)  F(12B)-C(12)-C(11)  117(1)  *superscripts refer to the symmetry operations: 1 x-1/2, 3/2-y, 1/2+z;^2 x-1/2, 1/2+y, z; 3  1+x, y, z;^  5  1/2+x, 1/2-y, z-1/2;^6 1/2+x, y - 1/2, z.  4  x-1, y, z;  about the Sn atom in the (CH3)2Sn(O2CCH3)2 and the RCH3) 2 Sn(02CCF3)2]„ proscribe their inclusion on the linear plot of isomer shift (6) vs. quadrupole coupling (AE Q). The MOssbauer spectrum of (CH 3 ) 2 Sn(O 2CH) 2 has been reported previously [24] and the isomer shift and quadrupole coupling are 1.45 and 4.72 mm•s -1 , respectively. These values fit reasonable well on the plot (Fig. 3-5) and show that the model proposed by Mallela  et al. [4] still holds for weakly nucluophilic anions. While the MOssbauer spectrum of (CH3) 2 Sn(O2CCH3)2 is not yet available, a spectrum of (CH3)2Sn(O2CCF3)2 has been obtained in the present study. The values obtained for the isomer shift and quadrupole  85 splitting obtained are 1.48 and 4.24 mnrs -1 , respectively. The data for this compound do not conform to the relationship between (5 and LIE Q . This is a reflection of the asymmetrical environment about the Sn atom and a departure of the C-Sn-C angle from linearity. Attempts to isolate (CH3)2Sn(O2CCH3)2 from solution resulted only in the isolation of the the "basic salt" [37] or the hydrolyzate, RCH 3 ) 2 Sn(0 2 CCH 3 )] 2 0. This confirms the extreme hydrolytic instability of (CH 3 ) 2 Sn(O 2 CCH3)2 as reported previously [30][31][36].  CONCLUSIONS The steric effects of the anions in dimethyltin(IV) carboxylate derivatives appear to have a great influence on the structure of these compounds. There does not appear to be a correlation between bonding modes and acidity in the derivatives studied in the present work. The overall structural chemistry of the carboxylates provides numerous examples for chelating as well as bridging (symmetric and asymmetric) groups. In contrast to the fluorosulfate group, which always functions as a bridging group (mainly symmetric bidentate), the carboxylate anions adopt various configurations, as is illustrated by the variety of structures encountered in this study. Of the dimethyltin(IV) carboxylates studied ((CH 3 ) 2 Sn(HC00) 2 , (CH 3 ) 2 Sn(CH 3 C00) 2 and (CH 3 )2Sn(CF 3 C00) 2) only the formate derivative adheres to the criteria required for the study of anion basicities by MOssbauer spectroscopy. The C-Sn-C bond angles in the acetate and the trifluoroacetate derivatives depart significantly from linearity, and it is not possible to use M6ssbauer spectroscopy as a probe for anion basicities in these two compounds. It is hoped that more structures become available for the stronger acid derivatives in order to further support the correlation between the two MOssbauer parameters and anion basicities.  86  REFERENCES  1. J. Barr, R.J. Gillespie and R.C. Thompson, lnorg. Chem., 3 (1964) 1149. 2. P.L. Fabre, J. Devyuk and B. Tremillon, Chem. Rev., 82 (1982) 591. 3. G. A. Olah, G.K. Surya Prakash and J. Sommer, "Superacids", Wiley, New York (1985). 4. S.P. Mallela, S. Yap, J.R. Sams and F. Aubke, Inorg. Chem., 25 (1986) 4327. 5. C.J. Evans and S. Karpel, "Organotin Compounds in Modern Technology", J. Organomet.  Chem. Libr. 16, Elsevier, Amsterdam (1985). 6. J.J. Zuckerman, R.P. Reisdorf and H.V. Elisse, in F.E. Brinckman and J.M. Bellamaed (eds.), "Organometallics and organometalloids, occurrence and fate in the environment", Am. Chem. Soc. Symp. Ser., 82 (1978) 388. 7. E.S. Hedges, Research, 13 (1960) 449. 8. A.G. Davies and P.J. Smith, in G. Wilkinson, F.G.A. Stone and E.W. Abel (eds.), "Comprehensive Organometallic Chemistry: The Synthesis, Reactions and Structures of Organometallic Compounds", Vol. 2., Pergammon, New York (1982) p. 519. 9. A.J. Crowe, P.J. Smith and G. Atassi, Chem. Biol. Interact., 32 (1980) 171. -  10. W.T. Piver, Environ. Health Perspect., 4 (1973) 61. 11. J.A.J. Thompson, M.G. Schaeffer, R.C. Pierce, Y.K. Chau, J.J. Cooney, W.R. Cullen and R.J. Maguire, "Organotin Compounds in the Aquatic Environment: Scientific Criteria for Assessing their Effects on Environmental Quality", NRC, Ottawa (1985). 12. F.A. Cotton and G. Wilkinson, "Advanced Inorganic Chemistry", 4th edn., Wiley, New York (1980) p. 403. 13. P.A. Yeats, Ph.D. Thesis, University of British Columbia, Vancouver, 1973. 14. T.H. Tan, J.R. Dalziel, P.A. Yeats, J.R. Sams, R.C. Thompson and F. Aubke, Can. J.  Chem., 50 (1972) 1843. 15. P.A. Yeats, B.F.E. Ford, J.R. Sams and F. Aubke, Chem. Comm., (1969) 791.  87  16. P.A. Yeats, J.R. Sams and F. Aubke, Inorg. Chem., 11 (1972) 2634. 17. E.O. Schlemper and W.C. Hamilton, Inorg. Chem., 5 (1966) 955. 18. F.H. Allen, J.A. Lerbscher and J. Trotter, J. Chem. Soc. A., (1971) 2507. 19. K. O'Sullivan, R.C. Thompson and J. Trotter, J. Chem. Soc. A., (1967) 2024. 20. A.G. Davies, H.J. Milledge, D.C. Puxley and P.J. Smith, J. Chem. Soc. A, (1970) 2862. 21. L.E. Smart and M. Webster, J. Chem. Soc. Dalton Trans., (1976) 1924. 22. J. Konnert, D. Britton and Y.M. Chow, Acta Crystallogr., Sect. B28, (1972) 180. 23. Y.M. Chow, Inorg. Chem., 9 (1970) 794. 24. Y. Sasaki, H. Imoto and 0. Nagano, Bull. Chem. Soc. Jpn., 57 (1984) 1417. 25. N.N. Greenwood, Chem. in Britain, 3 (1967) 56. 26. S.P. Mallela, S. Yap, J.R. Sams and F. Aubke, Rev. Chim. Miner., 23 (1986) 572. 27. L.E. Levchuk, J.R. Sams and F. Aubke, Inorg. Chem., 11 (1972) 43. 28. S. Karunanthy, F. Aubke, J.R. Sams, unpublished results 29. S.P. Mallela, S.T. Tomic, K. Lee, J.R. Sams and F. Aubke, Inorg. Chem., 25 (1986) 2939. 30. R. Okawara and E.G. Rochow, J. Am. Chem. Soc., 82 (1960) 3285. 31. Y. Maeda and R. Okawara, J. Organomet. Chem., 10 (1967) 247. 32. T.P. Lockhart, W.F. Manders and E.M. Holt, J. Am. Chem. Soc., 108 (1986) 6611. 33. Y. Maeda, C.R. Dillard and R. Okawara, Inorg. Nucl. Chem. Lett., 2 (1966) 197. 34. T.N. Mitchell, J. Organomet. Chem., 59 (1973) 189. 35. A.G. Maddock and R.H. Platt, J. Chem. Soc. A., (1971) 1191. 36. T.P. Lockhart, J.C. Calabrese and F. Davidson, Organometallics, 6 (1987) 2479. 37. F. Mistry, B.Sc. Thesis, University of British Columbia, Vancouver, 1987.  88  CHAPTER 4. FLUOROSULFATE DERIVATIVES OF GROUP 4 ELEMENTS  INTRODUCTION Among the very strong protonic or BrOnsted acids and superacids [1][2], the highest acidities observed are for "conjugate superacid systems" which usually consist of a strong protonic (BrOnsted) acid and a Lewis acid of equal strength [3]. Anhydrous hydrogen fluoride (HF) and fluorosulfuric acid (HSO 3 F) have emerged as the two strongest BrOnsted acids with identical Hammett (H0) values [4] of -15.1 reported for both of them [3][5]; both HF and HSO 3 F have found extensive use in many conjugate superacid systems [3]. On account of the wider liquid range of HSO3F, the greater ease of purification by distillation at atmospheric pressure [6], its compatibility with glass, and the large number of easily applicable physical techniques for characterization available, the use of HSO3F [7][8] offers clear advantages over the use of HF [9] in synthetic chemistry as well as for spectroscopic studies. The binary pentafluorides, particularly those of the elements in group 5 (Nb and Ta) and group 15 (As and Sb), have been frequently used as Lewis acids in conjugate superacid systems, together with either hydrogen fluoride or fluorosulfuric acid as the Br6nsted acids [3]. The conjugate superacid system HSO3F-SbF5, termed as "Magic Acid", is widely employed in the generation and stabilization of a wide range of electrophilic cations, even though facile fluoride vs. fluorosulfate exchange and concentration-dependent solute association cause several anions to be present in solution and introduce some complexity [9][10]-[12]. Some of this inherent complexity may be avoided in conjugate superacid systems of the type HSO 3 F-E(SO 3 F) n , where a high valent binary fluorosulfate acts as a Lewis acid and where, according to the general equilibrium:  2m HSO 3 F + E(SO 3 F),,=--"- m H 2 S0 3 F + + [E(SO3F) n+m ] m  -^-  1  -  89  stable fluorosulfate complexes are readily formed and isolated. In the past, four such conjugate superacid systems have been developed, albeit with some limitations. The use of the systems HSO 3 F-Au(SO 3 F) 3 [13][14] and HSO 3 F-Pt(SO3F) 4 [15], is limited by both the high price of the metals and the oxidizing ability of Au(III) and Pt(IV). In the case of the HSO 3 F-Au(SO 3 F) 3 system, the reduction of Au(III) by either gold powder or carbon monoxide has produced very interesting species such as Au 2+ (soh, ) [16][17], Au(CO)SO3F [18], and [Au(C0) 2] -/- [18][19]. In conjugate superacid systems, however, such redox reactions are undesirable since they degrade and consume the Lewis acid which results in a reduction of acidity. In the two superacid systems recently reported, HSO 3 F-Nb(SO 3 F)5 and HSO 3 FTa(SO 3 F) 5 [20] interference from redox reactions involving Nb(V) and Ta(V) is improbable; however, the two Lewis acids Nb(SO 3 F) 5 and Ta(SO 3 F) 5 are generated in situ and exhibit limited thermal stability. They both decompose via SO 3 elimination, resulting in the formation of fluoride fluorosulfates of Nb(V) and Ta(V) [20][21]. To increase the number of potential conjugate superacid systems in HSO 3 F, the syntheses of Lewis acids of the type M(SO3F)4 (M = Ti, Zr, or Hf) were carried out and are discussed in this chapter. In order to demonstrate the intrinsic fluorosulfate ion acceptor ability of these binary fluorosulfates, the syntheses, thermal stabilities and vibrational characteristics of their ternary fluorosulfate derivatives of the type [Cs 2 [M(SO 3 F) 6 ] are of considerable interest. According to a recent review [22], two general routes to binary and ternary fluorosulfates of the types M(SO3F)4 and Cs 2 [M(SO 3 F) 6 ] (M = Ti, Zr or Hf) may be contemplated: (i) ligand substitution of suitable MX4 or [MX 6 ] 2- precursors (X = halide, alkoxide or carboxylate), by HSO 3 F, its anhydride S 2 05 F2 or, in the case of X = Cl, by bis(fluorosulfuryl) peroxide or bromine(I) fluorosulfate.  90  (ii) metal oxidation using S 2 06 F2 as an oxidizer and HSO 3 F as the reaction medium either in the absence or presence of stoichiometric amounts of CsSO 3 F. Previous attempts to synthesize the M(SO3F) 4 (M = Ti, Zr) were confined to the use of route (i), and met with mixed success. Of the binary compounds, only the zirconium derivative has been reported so far. Zr(SO 3 F) 4 has been isolated in an analytically pure form by the solvolysis of zirconium tetrakis(trifluoroacetate), Zr(O2CCF 3 )4, in an excess of fluorosulfuric acid, HSO 3 F, and its infrared spectrum has been reported [23]. In most other instances, either incomplete substitution or complex decomposition of the initial products has been observed: e.g. the solvolyses of MC1 4 in HSO 3 F (M = Ti and Zr) studied by Hayek et al. [24] have reportedly yielded TiC12(SO3F) 2 and ZrF 3 SO 3 F, respectively. The aforementioned TiC1 2 (SO 3 F) 2 has also been obtained from the reaction of TiC1 4 with a slight excess of S20 6 F 2 at -20°C; when higher temperatures and a large excess of S 2 06 F2 were used, more extensive chloride substitution led to poorly-defined chloridecontaining materials [25]. In another attempt, the reaction of Ti(OCH 3 ) 4 with an excess of bis(fluorosulfuryl)oxide, S 2 0 5 F 2 , resulted in disubstitution only and Ti(OCH 3 ) 2 (SO 3 F) 2 has been isolated [26]. This chapter discusses the synthesis of the tetrakis(fluorosulfato) derivatives of the group 4 metals via oxidation with bis(fluorosulfuryl) peroxide in fluorosulfuric acid, and their characterization by vibrational spectroscopy. The syntheses and characterization of their ternary derivatives of the type Cs2[M(SO 3 F) 6 ] (M = Ti, Zr or Hf) are also discussed.  EXPERIMENTAL Reagents: Ti (-100 mesh, 99.9%), Zr (-80 mesh, 99.9%, packed under water) and Hf (-325 mesh, 99.6%, containing 2-3% Zr) were obtained from Morton Thiokol Inc., Alfa Products.  Caution: The finely divided forms of Group 4 metals are known to be pyrophoric. Although there were no explosions or violent reactions towards strong oxidizing reagents during the  91  course of this study, these metal powders should be used in small quantities, with suitable safeguards.  Synthetic reactions Except where mentioned, most of the reactions were carried out in 50 mL Pyrex round bottom flasks fitted with 4 mm Kontes Teflon stem stopcocks and standard taper B10 ground glass cones, as described in Chapter 2 (Fig. 2-2).  Ti(SO 3 F) 4 : To ca. 53 mg (1.106 mmol) of Ti metal, contained in the reaction vessel, about 5 mL of HSO 3 F were added in vacuo. After warming, the mixture was stirred at room temperature for ca. 15 min. After cooling the flask to -196°C, about 5 mL of S 2 06 F 2 were added to the mixture in vacuo. The resulting mixture was left stirring at room temperature overnight. After 24 hours, the mixture had a light green colour, and stirring was continued until the metal appeared to have been completely consumed (ca. 10 days). Upon opening the reactor at -196°C there was some residual pressure due to 0 2 (identified by mass spectrometry). The reactor was warmed to room temperature and as the product was pumped on, it became more viscous and eventually reached a constant weight (495 mg, 1.115 mmol). The residue was green in colour and had a resin-like consistency. Attempts to redissolve the product in HSO 3 F were unsuccessful.  Zr(S0 3 a4 : Prior to use, the Zr was dried in vacuo to remove all traces of water. In a typical reaction, 93 mg (1.020 mmol) of Zr metal were sealed in the reaction vessel, and about 5 mL of fluorosulfuric acid were added to this by vacuum transfer. The reactor was warmed briefly to about 100°C and the mixture was left stirring and allowed to cool to room temperature. After briefly pumping on the reactor, about 5 mL of S 2 06 F2 were added in vacuo. The reaction mixture was allowed to warm up to room temperature and left stirring overnight. The temperature was raised to 120°C and after 3 weeks all the Zr metal had been consumed to give  92  a white powder in suspension in the solution. The volatiles were pumped off at room temperature leaving behind 434 mg (0.890 mmol) of a white solid. During the reaction there was observable corrosion of the reactor and, hence, the weight balance must be regarded as unreliable. Analysis obtained (%): F = 16.4, S = 25.2, Zr = 17.69 (analysis expected (%): F = 15.64, S = 26.31, Zr = 18.71).  Hf(S0 3 E)4 : About 5 mL HSO 3 F were added by vacuum transfer to 235 mg (1.317 mmol) Hf metal. The mixture was heated briefly to about 100°C and allowed to cool to room temperature while stirring. About 5 mL of S 2 0 6 F 2 were vacuum transferred to the reactor and the mixture was allowed to warm up to room temperature overnight while stirring. The temperature was raised to 60°C for 24 hr, 100°C for one week, and then finally to 120°C for one week after which all the metal had been consumed. A white solid was suspended in the solution of HSO 3 F and S 20 6 F2. The reactor was cooled to -196°C and the oxygen was removed. The rest of the volatile products were pumped off at room temperature leaving behind a white solid weighing 804 mg (1.399 mmol). Analysis obtained (%): F = 13.9, S = 22.05, Hf = 31.94 (analysis expected (%): F = 13.22, S = 22.32, Hf = 31.06).  C_21 Ti(S03E)61: 53 mg (1.106 mmol) Ti and 513 mg (2.212 mmol) CsSO 3 F were weighed -  out in the reaction vessel. About 5 mL of HSO3F were transferred to this mixture in the drybox. About 5 mL of S 2 06 F2 were added to the solution via vacuum transfer and the reactor was allowed to warm to room temperature. Upon warming to room temperature there was some bubbling at the metal surface, and the bottom of the reactor got warm to the touch. The reaction mixture was left stirring at room temperature for about 5 days. The reactor was then evacuated at room temperature leaving behind a whitish-pale green solid weighing 979 mg (1.078 mmol). Analysis obtained (%): S = 21.35 (analysis expected (%): S = 21.18).  93  Cs 2 [Zr(S0 3 E161: 202 mg (2.214 mmol) of Zr and 1.032 g (4.449 mmol) of CsSO 3 F were weighed out in a 2 mm wall thickness pyrex ampoule 2 cm in diameter. The ampoule was connected to the vacuum line via an ampoule key as described previously [27]. About 5 mL each of HSO3 F and S 2 06 F2 were vacuum transferred to the ampoule and it was then sealed off. The mixture was left stirring at 110°C for 2 days upon which all the Zr was consumed. Upon letting the mixture stand at room temperature, two layers were observed. The lower layer converted, with time, to a waxy solid at room temperature, along with what appeared to be small crystals. Keeping the mixture at 4°C for about 10 weeks did not yield any crystals. The ampoule was cooled to -196°C and broken under vacuo. It was then evacuated overnight as it warmed to room temperature leaving behind 2.099 g (2.206 mmol) of a white solid. Analysis obtained (%): F = 12.1, S = 20.80, Zr = 9.82 (analysis expected (%): F = 11.98, S = 20.22, Zr = 9.59).  Cs2THf(S03E)61: 180 mg (1.008 mmol) of Hf and 369 mg (2.191 mmol) of CsCI were weighed out into a 50 mL pyrex round bottom flask fitted with a 4 mm Kontes Teflon stopcock and a B10 ground glass joint. About 5 mL HSO 3 F were vacuum transferred to the reactor and the mixture was allowed to warm up to room temperature while stirring. After 30 min, the reactor was pumped on to remove the HCl formed. About 5 mL of S 2 06 F2 were vacuum transferred to this mixture. The reactor was left stirring at 100°C for about 10 days during which time all the Hf was consumed. A white solid suspension in the HSO 3 F/S 2 06 F2 mixture was observed. The volatiles were pumped off leaving behind a white solid weighing 1.083 g (1.043 mmol). Analysis obtained (%): F = 10.9, S = 18.55, Hf = 19.86 (analysis expected (%): F = 10.97, S = 18.52, Hf = 17.18).  Solvolysis of ZrC1 4 in HSO 3 F: The solvolysis of ZrC1 4 was carried out at room temperature by dissolving ca. 135 mg of ZrC1 4 in HSO 3 F. The white product isolated after a short reaction time of ca. 1-2 days has the approximate composition ZrCl x (SO 3 F) 4 _, (x = 0.22).  94  When the reaction time was extended to three weeks, the resulting white product isolated after removal of all volatiles in vacuo was free of chloride. The analytical data (F = 17.4%, S = 20.63%) suggest a F:S ratio of 2:1 and best fit a composition of Zr00 5F1 5(SO3F)1 5• .  .  .  DISCUSSION The reported formation of ZrF 3 SO3 F in the reaction of ZrC1 4 with HSO 3 F at 100°C [24] is puzzling. In contrast to this observation, the reaction of ZrC1 4 with the related acid HSO3 CF 3 apparently results in complete substitution at 50°C to give Zr(SO3CF 3 ) 4 which is thermally stable up to 207°C [28]. Hayek et al. also report the formation of TiCl(SO 3 CF 3 ) 3 from TiC1 4 and HSO 3 CF3 under similar conditions [24]. A recent reinvestigation of this reaction under more forcing conditions (and with the anhydride S 2 05(CF3) 2 added to HSO3CF3) found further chloride substitution and the formation of a chloride-containing product of the composition TiC1 075 (SO 3 CF 3 ) 3.25 [29]. The same study also provides evidence for (NO2) 2 [Ti(03SCF3) 6 ], an anionic trifluoromethylsulfate complex of titanium, which is formed  via the reaction of Ti(NO 3 )4 with a mixture of HSO3CF3 and its anhydride S 2 05 (CF 3 ) 2 [29]. At this point, it should be mentioned that neither HSO 3 CF 3 nor its derivatives readily cleave off S0 3 . Before turning to the metal oxidation reactions, it appeared to be useful to reinvestigate the reaction between ZrC1 4 and HSO 3 F under milder conditions than those employed previously [24]. The white solid product obtained in our study from the solvolysis at room temperature has the approximate composition ZrCl x (SO 3 F) 4 _x (x = 0.22) and its vibrational spectrum is discussed below. It must be concluded that ZrC1 4 undergoes more extensive substitution of the chloride by fluorosulfate or trifluoromethylsulfonate than does TiC1 4 when HSO 3 F or HSO3CF3 are used [28][29]. The infrared spectrum of ZrClx (SO 3 F) 4 _x (x = 0.22), is identical to the spectrum of Zr(SO 3 F) 4 obtained by metal oxidation of Zr except for very slight differences in band shape and the occurrence of an additional band of medium intensity at 408 cm -1 , possibly due to a  95  Zr-Cl stretching vibration. The contention that chloride is almost completely replaced by fluorosulfate is also consistent with the rather low residual chlorine content of 1.6% as determined by microanalysis. The reaction time was extended to three weeks in order to achive complete substitution. The chloride-free product obtained had the approximate composition ZrClx (SO 3 F) 4 _x (x = 0.22). It must be concluded that decomposition of some of the fluorosulfate groups must have occurred via elimination of SO 3 and S 2 05 F 2 ; both are viable decomposition modes of metal fluorosulfates. The infrared and Raman specttra of this product are much simpler and appear to contain bands which are attributable only to a bidentate bridging fluorosulfate group. The infrared spectrum of the Zr(SO 3 F) 4 obtained by solvolysis of Zr(O 2 CCF 3 ) 4 in HSO 3 F reported by Singh et al. [23] is different and, particularly in the SO 3 stretching region where only three bands are reported, much simpler than the infrared spectra of all the metal fluorosulfates discussed below. While at first it appears likely that differences in instrumentation, spectral resolution, or sample preparation techniques may have contributed to these discrepancies, there is however another possible explanation. The product obtained in our study from the solvolysis of ZrC1 4 in HSO 3 F after a reaction times of 3 week shows a strong spectral similarity to the one reported by Singh et al. [23], and all vibrations in the SO 3 F-stretching range are assignable to a bidentate, possibly bridging SO3F-group in both instances. It hence appears that the Zr(SO 3 F) 4 that was initially formed became degraded in much the same manner as the solvolysis product of ZrC1 4 in HSO3 F obtained in this study. Metal oxidation by bis(fluorosulfuryl) peroxide (S2O 6 F2) may be carried out at reaction temperatures of up to 150°C to yield either binary or ternary fluorosulfates according to the following reactions:  HSO3 F M + n/2 S 2 06 F2 ^ > M(SO3F) n^-2-  96  HSO 3 F ^ 2 CsSO 3 F + M + 4 S206 F 2 ^ > Cs2[M(SO3F)6] -3-  This approach has been used previously with a number of metals (Pt [15], Nb, Ta [20], Pt, Pd [30], Ir [31], Ru [32], Sn and Ge [33]) all of which form ternary Cs2[M(SO 3 F) 6 ] as well. Details on the oxidation of Ti, Zr and Hf by S 2 06 F2 in HSO 3 F are summarized in Table 4.1, where the results are compared to some previous studies [15][30]-[33] with respect to reaction times, reaction temperatures and the thermal stability of the resulting products. As can be seen, all three binary fluorosulfates of the group 4 elements have been obtained along with their cesium salts, and all the compounds, with the exception of Ti(SO 3 F) 4 , are hygroscopic white solids. The seemingly long reaction times required to oxidize Zr and Hf are, in part, due to the careful approach taken during the initial stages of the reactions on account of the well-documented pyrophoric and easily oxidizable nature of the metal powders. In all our reactions, however, the metal powders reacted sluggishly and complete oxidation required extended reaction times and relatively high temperatures. In some instances 02 was produced as a byproduct, possibly as a result of the reaction of SO 3 F radicals with the wall of .  the reactor at elevated temperatures, as reported previously [34]. The formation of 02 together with bis(fluorosulfuryl) oxide, S 2 05 F 2 , has also been reported during the oxidation of iridium [31] and rhodium [35] where prolonged reaction times and similarly high reaction temperatures of 140-150°C were employed. Two apparent inconsistencies during the syntheses of Zr(SO 3 F) 4 and Hf(SO3F)4 require further explanation. The final weight of Zr(SO 3 F) 4 appears to be slightly less than expected. This can be rationalized in terms of the corrosion of the inner walls of the glass reactor during the prolonged reaction times; this hypothesis is also consistent with the formation of SiF 4 (detected by infrared spectroscopy) during the reaction. On the other hand, the weight of Hf(SO 3 F) 4 obtained appears to be slightly higher, suggesting formation of 1.399 mmol of the  97 Table 4.1: Binary and Ternary Metal Fluorosulfate Derivatives obtained by Oxidation with S 2 06 F2 in HSO3 F Compound  ^  Reaction^Reaction Thermal^Description^Ref.' Temp (°C) Time (d)^Stability  Ti(SO 3 F) 4^25-60^10^ yellow-green resin^tw Zr(SO 3 F) 4^25-120^21^up to 180°C^white solid^tw Hf(S 03F)4 b^25-120^21^up to 190°C^white solid^tw Sn(SO3F)4^25^0.5^up to 196°C^white solid^33 Pt(SO 3 F) 4^120^2^220°C*^yellow-orange solid^15 Ir(SO 3 F) 4e^60-140^6.5^up to 150°C^dark brown solid^31 Pd(SO 3 F) 3 d^25-120^3^up to 180°C^black brown solid^30 Ru(SO 3 F) 3^60^ca. 1^up to 140°C^red brown solid^32 GeF 2 (SO 3 F) 2e^50^3^146°C*^white solid^33 Cs 2 [Ti(SO 3 F) 6 ] 25^5^60-65°C*^white solid^tw Cs 2 [Zr(SO 3 F) 6 ] 25+110^1+2^260-265°C*^white solid^tw Cs2 [Hf(SO 3 F) 6 ] 100^10^262°C*^white solid^tw Cs 2 [Sn(SO 3 F) 6 ] 25^0.5^249-253°C*^white solid^33 Cs 2 [Pt(SO3F) 6 ]^80^3^260°C*^light yellow solid^15 Cs2 [Ir(S 03F)6]^150^10^ca. 150°C *^pale orange solid^31 C S2[Pd(S 03 F)6] 120^3^up to 200°C^dark red solid^30 Cs2 [Ru(SO 3 F) 6 ] 60^1^not stated^orange solid^32 Cs2 [Ge(SO 3 F) 6 ] 50^2^242°C*^white solid^33 * decomposition point; 'tw = this work; b contains 2-3% Zr(SO 3 F) 4 ; obtained by the oxidation of Ir(SO 3 F) 3 with S20 6 F2 [31]; d formulated as Pd(II)[Pd(IV)(SO 3 F) 6 ] [30]; e obtained in 75-80% yield when Ge is oxidized by S 2 06 F 2 in HSO3F [33].  98  product rather than the expected 1.317 mmol. As stated in the Experimental (vide supra), the hafnium powder employed contains 2-3% Zr according to the supplier; this would produce 23% Zr(SO 3 F) 4 in the final product which accounts for the observed weight difference reasonably well. Both Zr(SO 3 F) 4 and Hf(SO 3 F) 4 are white hygroscopic powders of comparable thermal stabilities, and both compounds give nearly identical infrared spectra (vide  infra). It is therefore not possible to detect the presence of a small amount of zirconium(IV) fluorosulfate in hafnium(IV) fluorosulfate by vibrational spectroscopy. The appearance and thermal stability of Zr(SO 3 F) 4 agree reasonably well with the Zr(SO 3 F) 4 product obtained from Zr(O 2 CCF 3 ) 4 and HSO 3 F [23]. In all the cases summarized in Table 4.1 and those discussed elsewhere [22], binary fluorosulfates are found to be relatively high-melting, hygroscopic solids. The appearance of Ti(SO 3 F) 4 is, however, unprecedented. The compound Ti(SO 3 F) 4 is a waxy, resin-like material of low thermal stability which does not redissolve in HSO 3 F. This does not allow either sublimation or recrystallisation from HSO 3 F. In addition, a quantitative analysis was not possible because reliable weights of the sample for analysis were not obtainable. Also, attempts at microanalysis resulted in a rapid corrosion and shattering of the quartz tube in the apparatus. The yellow-greenish colour of the product is not surprising; all titanium(IV) chlorosulfates are described as yellow solids [24][25][28][29]. Unlike Zr(SO 3 F) 4 or Hf(SO 3 F) 4 , all the samples of Ti(SO 3 F) 4 studied by infrared spectroscopy vigorously attacked silver halide windows. There is some evidence that Ti(SO 3 F) 4 obtained in the manner described above may not be pure in spite of the good weight balance and microanalysis. Spectra obtained on samples pressed between silicon discs show a medium-weak broad absorption at approximately 3100 cm -1 , which suggests the retention of residual amounts of fluorosulfuric acid in the sample. All attempts to remove the acid by heating in vacuo resulted in some sample decomposition and subsequent weight loss. The presence of HSO 3 F may be responsible for the difference in appearance of Ti(SO 3 F) 4 . This situation is reminiscent of the oxidation of Nb or Ta by S20 6 F 2 in HSO3 F, which leads to the  99  in situ formation of M(SO 3 F) 5 (M = Nb or Ta) which are too labile to permit isolation from solution [20]. The infrared spectrum of Ti(SO 3 F) 4 shows SO 3 stretches at 1390(vs), 1251(ms), 1110(w) 1050(vs), 960(vs,sh) and 880 in the stretching range with bands at 640(s), 585(ms), 551(m) and 450(s) caused by deformation vibrations; all the bands are similar to those observed in the infrared spectra of Zr(SO3F) 4 and Hf(SO3F) 4 (vide infra). The infrared spectra of Zr(SO3F) 4 and Hf(SO3F) 4 between 1500 and 350 cm -1 are shown in Fig. 4-1. The most intriguing observation is the very close similarity in band positions, intensities and shapes for both the compounds. Only in the low-wavenumber region (below 500 cm -1 ) are small discrepancies visible. The identical oxidation states, atomic and covalent radii for Zr and Hf [36], which are responsible for many of the identical chemical and structural features of their compounds, are thought to be the reason for the observed spectral similarities. This observation also strongly suggests a common molecular structure for both compounds. The band positions and estimated intensities for both compounds are listed in Table 4.2 and compared to the previously reported infrared spectra of [Au(SO 3 F)3] 2 [14][37], and the tetrakis(fluorosulfates) of Pt [15], Ir [30] and Sn [33]. The previously published spectra of Zr(SO3F)4 [23] and those of the products of the solvolysis of ZrC14 in an excess of HSO 3 F either after short or long exposure times are also included in this comparison. The presence of only a bidentate bridging fluorosulfate group in some tin(IV) derivatives has been confirmed for (CH 3 ) 2 Sn(SO 3 F) 2 [38] by X-ray diffraction and has been suggested for SnC12(SO3F)2 [33] and SnF2(SO3F)2 [39] based on their  119 Sn MOssbauer  spectra; this observation aids in the approximate band assignments in Table 4.2. Three bands at different positions are expected in the S0 3 -stretching region for the two coordination modes of the fluorosulfate group with C s (or even lower) local symmetry. Monodentate, terminal fluorosulfate groups, are characterized by three bands that are usually found at 1500-1420, 1270-1220 and 1000-850 cm -1 , while bidentate groups are best recognized by bands at 1420-  420 470 z 358 350 / 564 (A  Hf(S0,94  430  Zr(SO3 F  t 46 5  )  642 /^ ),  385 475^380  4  633  570 600  1418^  1073  1393 1380 1150^  1 1110  Wavenumber(cni 1 ) Fig. 4-1. Infrared Spectra of Hf(SO 3 F)4 and Zr(SO 3 F)4  Table 4.2. Infrared Spectra of [Au(SO,F),], [13][14] and the Tetrakis(fluorosulfates) of Zirconium, Hafnium, Tin [44], Platinum [15] and Related Compounds [23]. Sn(SO 3 F)4 V[cm-I ] Int  Pt(S0,0 4 Zr(S0304[23 ] Zr20F3(SO30, Wee] Int 9[cm -1 ] Int V[cm - 1 Int  1415 s,sh 1390 s 1378 vs  1438 s 1411 s 1395 w,sh  1440  vs,b  1274 s 1170 w,sh 1150 vs  1269 s 1180 w 1150 vs  1232  1225  s  1130 s  1150  vs  1116  s  1140 m,b  1110 s,sh 1073 m,sh  1120 s,sh 1078 m,sh  1085 s  1070  vs  1068  ms  1060 s  1022 m 972 w,sh  1018 m,w 970 vw,sh  1000  vs  [Au(SO,F) 3 ] 2 V[cm- 1  Zr(S0,94 V[cm-1] Int  1442 vs 1425 s  1418 s,sh 1393 vs 1380 vs  1240 s 1220 s,sh 1135 ms 1055  HRS0304 V[cml Int  1403  s  1395 m  Approximate Description vS03asym for bridging and terminal SO,F-group vSO3 sym, terminal SO,F group vS03 asym, bridging SO,F group } vS0 3 sym bridging  s  960 s,b 920 s,sh 895 s,b 820 s  862 853  682 s, 670 s,sh 610w, 590s 582s,sh, 550s 460 m  } vSO-M terminal SO,F groups  980 w,sh vs,b  920 m 850 m 832 s  900  860 s,sh 853 s  642 w,sh 633 m  640 w 628 m  657 w 628 m  670 640  s,b s,b  658  600 s 570 s  594 s 564 s  585 m 550 s  580 540  s m  593 563  475 465 430 385 380  470 460 420 358 335  460 w 438 m 427  474 430  br  328 m  320  w  s,sh s  s s,sh vw m m,sh  m m,sh vw m m  862  m 835 m  s  1^vSF for bridging and terminal groups  650 m  } vM-O + deformation mode  590 m 550 m  ) OS0 3 bending vibrations  440 w 295  s  various deformation modes for rocking vibrations and M-0 skeletal vibrations  0  102  1350, 1170-1100 and 1080-1060 cm -1 . A similarly clear differentiation between the deformation modes is not possible on account of the rather narrow wavenumber range (600400 cm -1 ) and the extensive band overlap in this area. The infrared spectra summarized in Table 4.2 provide strong evidence for the presence of fluorosulfate bridged polymers or oligomers. Oligomeric or polymeric structures are expected for all the M(SO 3 F) 4 compounds, and this is consistent with their limited solubilities in HSO 3 F. This aggregation is also expected to result in complex vibrational spectra. The presence of two different types of fluorosulfate groups, monodentate and bidentate bridging, has been confirmed for the dimeric gold(III) fluorosulfate by an X-ray diffraction study [37] and is also suspected for the remaining fluorosulfates listed in Table 4.2 based on the observed spectral similarities. There are no obvious reasons why three of the metal fluorosulfates, [Au(SO 3 F) 3 ] 2 [14], Pt(SO 3 F) 4 [15] and Ir(SO 3 F) 4 [30], two of which act as Lewis acids in conjugate superacid systems [14][15], are quite soluble in HSO 3 F while the rest of the fluorosulfates are virtually insoluble. In all instances, unfortunately, the binary fluorosulfate derivatives of the group 4 elements exhibit insufficient solubility in fluorosulfuric acid. This precludes conductivity, NMR or Hammett acidity studies on the tetrakis(fluorosulfates) of the group 4 elements in this solvent. It appears that bridging by the fluorosulfate groups leads to polymers similar to those suggested for Sn(SO 3 F) 4 by 119 Sn M6ssbauer spectroscopy [33]. The oxidation of the group 4 metals by S 2 06 F 2 in HSO 3 F in the presence of CsSO 3 F was carried out in order to study the fluorosulfate ion acceptor ability of the binary fluorosulfates of Ti, Zr and Hf. There is also the potential of obtaining crystalline salts of the type Cs 2 [M(SO 3 F) 6 ] (M = Ti, Zr or Hf) for single crystal X-ray diffraction studies. Cesium fluorosulfate was chosen in all instances because an incomplete reaction resulting in residual CsSO 3 F can be readily recognized by the occurrence of an infrared band at 728 cm -1 , assigned as v2 and attributed to a sulfur-fluorine stretching vibration [32]. For  103  the fluorosulfates of other group 1 metals, vS-F is found at gradually increasing wave numbers [40], which makes an unambiguous assignment more difficult. In addition, the vibrational spectra of Cs 2 [M(SO 3 F) 6 ] salts are less complex, because of the larger, spherical shape of Cs + and its low ionic potential, than those of other group 1 cations or of the hetero cations like C10 2 + or NO [15][30]-[33] reported occasionally. As observed for all the reactions summarized in Table 4.1, shorter reaction times and lower reaction temperatures may be used when CsSO 3 F is present in the reaction mixture. The occasional observation of two liquid layers during the initial stages of the reaction has also been reported in a recent study of the HSO 3 F-S 20 6 F 2 system [41]. The addition of CsSO 3 F or another alkali metal fluorosulfate reduces the acidity of HSO 3 F and limits the solubility of S 2 06 F 2 , which has been found to behave as a weak base in HSO 3 F [41]. The high thermal stabilities of the ternary fluorosulfates are noteworthy; the compounds Cs 2 [M(SO 3 F)6] where M = Zr, Hf, Ge [33], Sn [33] or Pt [15] all have decomposition points around 250°C. The high thermal stability of all the anionic complexes may also be considered an indication of the acceptor ability of the binary fluorosulfates; the M(SO 3 F) 4 compounds have not been isolated for M = Pd [30], Ru [15] and Ge [33], although their ^ [Pd Iv (SO3F)6] [30], Ru(SO 3 F) 3 [15], and tris(fluorosulfates), Pd(SO 3 F) 3 and GeF 2 (SO 3 F) 2 [33] have been obtained. There must, however, be some very limited solubility of all metal fluorosulfates in HSO 3 F when CsSO 3 F is present in order to allow the formation of ternary fluorosulfates of the type Cs 2 [M(SO 3 F)6] for M = Ti, Zr or Hf. The vibrational spectra of the ternary fluorosulfates are summarized in Table 4.3 and are compared to the Raman spectrum of Cs2 [Sn(SO 3 F) 6 ] [33], for which octahedral coordination for tin is indicated by the 119 Sn MOssbauer spectrum. The similarity in the vibrational spectra (vide infra) of the ternary fluorosulfate derivatives of titanium, zirconium and hafnium with one another (See Fig. 4-2) and with those of other Cs 2 [M(SO3F)6] compounds reported previously indicates that these compounds have very similar structures. The principal SO 3 F-stretching vibrations give rise to  _ Zr Hf and Sn 33b_ . = Ti,^ Approximate description of bands  M=Ti; Ra Av[cm'] hit  M=Zr; IR V [cml Int  M=Hf; IR ii[cml Int  M=Hf; Ra Av[cm-1 ] Int  M=Sn; Ra tiv[cml Int  1400 vs,b  1405 w,b  1420 m,sh 1359 vs,b  1390 s,sh 1370 vs  1398 w 1382 m  1407 m 1399 m  } vS0, asym  1276 w 1211 vs  1278 vs 1210 w  1215 vs,b 1204 s,sh  1265 w 1218 vs  1280 s 1215 m  1270 s 1218 w  }  1106 m  1091 s 995 s,sh  M=Ti; IR N/ [cml int  vS0 2. sym  990 ms,b 970 m,sh  1109 m,s 1077 w,sh 1050 vs —1020 m,sh  1150 vw 1110 m 1050 s,sh 1015 vs  815 s 790 s,sh 700 vw  813 m,b  844 ms 790 s  810 vs 795 s,sh  858 s 770 w  828 m 811 m  651 mw 619 w  625 m  620 m,s  620 m  626 m 620 m,sh  625 s  vM-O + def. mode  —580 m,sh 566 s  580 vw 563 ms  585 m,sh 567 s  575 s,sh 565 m  570 m 555 m  578 ms 560 ms  OSO, asym 8S0, sym combination bands  415 m  431 m 418,407 w 345 w  1110 vw 980 vs,b  520 vw 430 ms 420 m,sh  vS0-M  1028 s }  vSF  495 vw,sh 423 ms,b  425 ms  405 s 259 s 170 w  405 w 240 ms  260 m  M-0 skeletal vibrations pSO3F M-0 skeletal vibrations + deformation modes  520 430  405 495  1370  1050 ^1015 1218 Wavenumber (cm')  Fig. 4-2. Infrared Spectra of Cs 2 [Ti(SO 3 F) 6] and Cs2[Hf(SO3F)6]  106  bands at 1400, 1280-1210, 1110-970 (vS-O) and 860-770 (vS-F), which are indicative of the presence of anionic, monodentate SO 3 F groups in the various [M(SO 3 F) 6 ] anions. Some band proliferation, probably caused by vibrational coupling, frequently results in the formation of shoulders, particularly in the infrared spectra. Very similar vibrational spectra are reported for all the other Cs2[M(SO3F) 6] salts as listed in Table 4.3, and their vibrational assignments have been adopted from Ref. 33. Both the infrared and the Raman spectra obtained for Cs2 [Ti(SO 3 F) 6 ] are in complete agreement with the spectra listed in Table 4.3 and those published previously [15][30]-[33]. The titanium derivative is, however, less thermally stable than the other compounds listed in Table 4.3. For M = Zr, several attempts were also made to obtain the anions [Zr(SO3F)71 3- or [Zr(SO3F)8] 4- by using a CsSO3F:Zr stoichiometric ratio of 3:1 or 4:1; in all instances, however, the infrared spectra of the resulting solid products indicated the presence of uncomplexed CsSO3F in addition to Cs2 [Zr(SO 3 F) 6 ]. It appears that the octahedral coordination around the zirconium cannot be expanded to give a 7- or 8-coordinated zirconium atom.  CONCLUSIONS Metal oxidation by S 2 06 F2 in HSO 3 F provides a simple, versatile synthetic route to binary, as well as ternary, fluorosulfates wherever the resulting reaction products have sufficient thermal stabilities in excess of the reaction temperatures needed for their synthesis, as indicated by the examples in Table 4.1. The oxidation of the group 4 metals (Ti, Zr and HO by S206 F2 in HSO3 F, at elevated temperatures (25-120°C) and after prolonged reaction times (10 to 21 days), leads to the corresponding fluorosulfates having the composition M(SO 3 F) 4 , M = Ti, Zr or Hf. The zirconium and hafnium compounds are extremely similar in their appearance (hygroscopic white solids, thermally stable to ca. 180°C) and give identical infrared spectra. The titanium derivative, Ti(SO 3 F) 4 , is a greenish-yellow resin-like material which retains small amounts of HSO3F.  107  All three binary fluorosulfates appear to be virtually insoluble in fluorosulfuric acid, which limits their use in solution studies such as electrical conductivity and the measurement of Ho values in Hammett acidity functions [5]. However, the intrinsic acceptor ability towards SO 3 F ions has been demonstrated for all three materials by the isolation of ternary -  fluorosulfates of the composition Cs2[M(SO3F)6], where M = Ti, Zr or Hf, and the relatively high thermal stability (to ca. 260°C). The ternary derivatives appear to be isostructural to other Cs 2 [M(SO 3 F) 6 ] salts where M = Ru, Ir, Pd, Pt, Ge and Sn, as is evident from Table 4.3. It should be remembered that the intrinsic acceptor ability of the metal centre may be manifested in two ways: (i) the propensity to function as Lewis acids towards weakly basic donor ligands like the SO 3 F ion; and (ii) the ability to form polymers with polydentate -  bridging anionic groups. This ability is expected to be strengthened by anionic groups that show a strong tendency to function as polydentate ligands, e.g. the fluorosulfate group. It is possible to use binary fluorosulfates in conjugate HSO3F superacid systems only where there is a limited tendency towards polymer formation: e.g. Au(SO3F)3, which is dimeric in the solid state [37] and possibly also in HSO 3 F solution [42]. It is interesting to note that the two binary fluorosulfates that do meet these criteria and are strong Lewis acids in HSO 3 F, Au(SO 3 F) 3 [14] and Pt(SO 3 F) 4 [15], were first synthesized by Professor Cady and his students over 20 years ago [43].  108  REFERENCES  1. R.J. Gillespie and T.E. Peel, Adv. Phys. Org. Chem., 9, (1972) 1. 2. R.J. Gillespie and T.E. Peel, J. Am. Chem. Soc., 95, (1973) 5173. 3. G.A. Olah, G.K.S. Prakash and J. Sommer (eds.) "Superacids", John Wiley and Sons, New York (1985). 4. L.P. Hammett and A.J. Deyrup, J. Am. Chem. Soc., 54, (1932) 2721. 5. R.J. Gillespie and J. Liang, J. Am. Chem. Soc., 110, (1988) 6053. 6. J. Barr, R.J. Gillespie and R.C. Thompson, Inorg. Chem., 3, (1964) 1149. 7. R.C. Thompson, in G. Nickless (ed.) "Inorganic Sulphur Compounds", Elsevier, Amsterdam (1968) p. 587. 8. A.W. Jache, Adv. Inorg. Chem. Radiochem., 16, (1974) 177. 9. H.H. Hyman and J.J. Katz, in T.C. Waddington (ed.) "Non-Aqueous Solvent-Systems", Academic Press, London (1965) p. 47. 10. R.C. Thompson, J. Barr, R.J. Gillespie, J.B. Milne and R.A. Rothenbury, Inorg. Chem.,  4, (1965) 1641. 11. G.A. Olah and A. Commeyras, J. Am. Chem. Soc., 91, (1969) 2929. 12. P.A.W. Dean and R.J. Gillespie, J. Am. Chem. Soc., 92, (1970) 2362. 13. K.C. Lee and F. Aubke, Inorg. Chem., 18, (1979) 389. 14. K.C. Lee and F. Aubke, Inorg. Chem., 19, (1980) 119. 15. K.C. Lee and F. Aubke, Inorg. Chem., 23, (1984) 2124. 16. H. Willner, F. Mistry, G. Hwang, F.G. Herring, M.S.R. Cader and F. Aubke, J.  Fluorine Chem., 52, (1991) 13. 17. F.G. Herring, G. Hwang, K.C. Lee, F. Mistry, P.S. Phillips, H. Willner and F. Aubke,  J. Am. Chem. Soc., 114, (1992) 1271. 18. H. Willner and F. Aubke, Inorg. Chem., 29, (1990) 2195.  109 19. H. Willner, J. Schaebs, G. Hwang, F. Mistry, R. Jones, J. Trotter and F. Aubke, J. Am. Chem. Soc., 114, (1992) 8972.  20. W.V. Cicha and F. Aubke, J. Am. Chem. Soc., 111, (1989) 4328. 21. D. Zhang and F. Aubke, J. Fluorine Chem., 58, (1992) 81. 22. F. Aubke, M.S.R. Cader and F. Mistry, in G.A. Olah, R.D. Chambers and G.K.S. Prakash (eds.) "Synthetic Fluorine Chemistry", John Wiley-Interscience, New York (1992) p. 43. 23. S. Singh, M. Bedi and R.D. Verma, J. Fluorine Chem., 20, (1982) 107. 24. E. Hayek, J. Puschmann and A. Czaloun Monatsch., 85, (1954) 359. 25. J.R. Dalziel, R.D. Klett, P.A. Yeats and F. Aubke, Can. J. Chem., 52, (1974) 231. 26. D.G. Niyogi, S. Singh and R.D. Verma, Can. J. Chem., 67, (1989) 1895. 27. W. Gombler and H. Willner, J. Phys. E. Sci. Instrum., 20, (1967) 1286. 28. M. Schmeisser, P. Sartori and B. Lippsmeier, Chem. Ber., 103, (1970) 868. 29. H. Moulayel Mustapha and J.P. Pascal, J. Fluorine Chem., 55, (1991) 63. 30. a) K.C. Lee and F. Aubke, Can. .1. Chem., 55, (1977) 2473. b) K.C. Lee and F. Aubke, Can. J. Chem., 57, (1979) 2058. 31. K.C. Lee and F. Aubke, J. Fluorine Chem., 19, (1982) 501. 32. P.C. Leung and F. Aubke, Can. J. Chem., 62, (1985) 2892. 33. a) S.P. Mallela, K.C. Lee and F. Aubke, Inorg. Chem., 23, (1984) 653. b) P.A. Yeats, J.R. Sams and F. Aubke, Inorg. Chem., 12, (1973) 328. 34. F.B. Dudley and G.H. Cady, J. Am. Chem. Soc., 85, (1963) 3375. 35. P.C. Leung, G.B. Wong and F. Aubke, J. Fluorine Chem., 35, (1987) 607. 36. L. Pauling, "The Nature of the Chemical Bond", Cornell University Press, Ithaca, New York (1960). 37. a) H. Willner, S.J. Rettig, J. Trotter and F. Aubke, Can. J. Chem., 69 (1991) 391. b) K.C. Lee, Ph.D. Thesis, The University of British Columbia, 1980. 38. F.A. Allen, J. Lerbscher and J. Trotter, J. Chem. Soc. A, (1971) 2507.  110 39. L.E. Levchuk, J.R. Sams and F. Aubke, Inorg. Chem., 11 (1972) 43. l Chem., 372, (1970) 40. A. Ruoff, J.B. Milne, G. Kaufmann and M. Leroy, Z. Anorg. Aug.  119. 41. W.V. Cicha and F. Aubke, Can. J. Chem., 68, (1990) 102. 42. W.V. Cicha, K.C. Lee and F. Aubke, J. Solution Chem., 19 (1990) 609. 43. W.M. Johnson, R. Dev and G.H. Cady, Inorg. Chem., 11 (1972) 2260 . 44. P.A. Yeats, B.L. Poh, B.F.E. Ford, J.R. Sams and F. Aubke, J. Chem. Soc. A, (1970) 2188.  111  CHAPTER 5. FLUOROSULFATE DERIVATIVES OF MOLYBDENUM  INTRODUCTION Molybdenum and tungsten (both group 6 elements) are known to be easily oxidizable to the hexavalent state even when relatively weak oxidizing agents are used. Neutral halogen derivatives of molybdenum(VI) and tungsten(VI) are found in the form of binary halides (MX 6 ; X = F, CO and oxyhalides (MOX 4 and MO2 X2 ; X = F, Cl). In the case of the molybdenum species, most of the oxyhalides are octahedrally coordinated about the central atom via oxygen bridging, whereas MoOF 4 appears to have bridging fluorine [1]-[6]. Substitution of a fluorine group with a fluorosulfate group introduces an additional potentially bridging group and reduces the propensity of oxygen to act as a bridging ligand. It has long been known that bis(fluorosulfuryl) peroxide, S 2 O6 F2 , can be used in the synthesis of binary and ternary fluorosulfates [7]. There are, however, no binary fluorosulfate derivatives of molybdenum reported in the literature [8]. A previous report [9] of the reaction of either molybdenum metal or molybdenum hexacarbonyl, Mo(CO)6, with bis(fluorosulfuryl) peroxide, S2O6F2, has suggested the formation of two products — a white solid, MoO2 (SO 3 F) 2 , and an orange liquid, thought to be MoO(SO 3 F) 2 . The existence of MoO(SO 3 F) 2 , with molybdenum in a +4 oxidation state, in the presence of a strong oxidizing agent such as S 2 O 6 F2 is unexpected [9], and a reinvestigation of the reaction of Mo with S2O6F2 by Lee [10] has shown the orange liquid to be MoO(SO 3 F) 4 . Neither of the two oxofluorosulfato derivatives was found to be a SO 3 F - acceptor, and neither one gave rise to the formation of anionic complexes. The vibrational data for MoO(SO 3 F) 4 are seen to be in extremely good correspondence with those available for the analogous tungsten compound, WO(SO 3 F) 4 [11], which was reported to be the only compound obtained in the reaction of S 2 O 6 F 2 with tungsten at  ca. 100°C [12]. The sharp Raman band at 933 cm -1 is attributed to the Mo=O stretching vibration and a terminal Mo=O bond is suggested. The SO 3 F band positions indicate the  112  presence of both bridging and terminal SO3F groups and there appears to be an absence of strong vibrational coupling, evidenced by the lack of band proliferation. The vibrational spectra for Mo0 2 (SO 3 F) 2 show the M(0)2 stretches as a pair of sharp  peaks at ca. 994 and 955 cm -1 in both the infrared and the Raman spectra. They compare well with the terminal M(0)2 stretches reported for MoO2C12 [4]. The relatively high intensity of the vsy . at 994 cm -1 in the infrared indicates a non-linear geometry for the M(0)2 moiety. The characteristic vas SO2 at ca. 1170 cm -1 in the infrared suggests the presence of bridging fluorosulfate groups, and the vS-F at 873-875 cm -1 points to a strong interaction between the molybdenum atom and the fluorosulfate group [13]. The SO 3 F bands correspond well with those observed for SnC12(SO3F)2 [14], even though a linear CI-Sn-C1 grouping appears to be present for SnC12(SO3F)2, as suggested by both  119 Sn MOssbauer and vibrational  spectroscopy. Thus, the geometry about the Mo(VI) centre is proposed to be octahedral with  cis terminal oxo groups on one side and bridging SO 3 F groups on the other side of a distorted octahedron. The Mo02 (SO 3 F) 2 species has also been obtained by the thermal decomposition of MoO(SO 3 F) 4 at ca. 100°C. One may postulate the following scheme for the formation of the two oxo-fluorosulfato derivatives of molybdenum(VI): Mo + 3 S 2 06 F 2  > "M o(SO 3 F) 6 "^ - 1 -  "Mo(SO 3 F) 6 "  > MoO(SO3F)4 + S2O5F2  -2-  MoO(SO 3 F) 4  > M 0 02(S03 9 2 + S2O5F2  -3-  This reaction sequence is supported by the observation of S 2 O5 F 2 in the volatile products. However, the proposed initial intermediate binary fluorosulfate, Mo(SO 3 F) 6 , has neither been observed nor isolated. For many elements, in addition to the oxyfluorosulfates, there have also been reports on several fluoride-fluorosulfates. As discussed in Chapter 1, fluoride-fluorosulfate derivatives are also known to play a useful role in superacid systems based on HSO3F, and possibly also HF. The oxidative addition of S 2 06 F2 to lower valent element fluorides has  113  been used in the past to synthesize fluoro-fluorosulfato derivatives of the type EF 3 (SO3 F) 2 (E = As [15] and Sb [16]) and EF 3 SO3 F (E = Ce [17]). Some of the volatile hexafluorides, EF6 , have seen some use in isotope separation either by mass diffusion in a hydrogen stream or by using laser excitation. For a hexafluoride, EF6, only five vibrational modes are active for the purpose of vibrational excitation to enhance reactivity, and only two of these are infrared active. A reduction in symmetry from Oh would produce more accessible vibrations for laser excitation. The synthesis of volatile derivatives of the type EF 5 SO3 F (E = S, Mo and U) is also, therefore, seen to be of importance in the field of isotope separation by laser excitation and diffusion. The oxidative addition of bis(fluorosulfuryl) peroxide, S 2 06 F2 , to low valent fluorides has been used previously to synthesize AsF3(SO3F) 2 [15], and SbF 3 (SO 3 F) 2 [16]. These compounds are viscous oily materials strongly associated via SO3F bridges with the fluorines present as terminal ligands. An analogous reaction of S 2 06 F2 with MoF5 was carried out in the present study with the intention of the synthesizing MoF5SO 3 F according to:  2 MoF 5 + S 2 06 F2  > 2 MoF5SO 3 F^ -4-  There are several tetra-A-carboxylato derivatives of Mo 2 4+ in which the carboxylate ligands may be substituted by various ligands [18]-[21]. Reports indicate the presence of a strong quadruple bond between the two molybdenum nucleii [19][22][23], involving a-, 2x w-, and 6-bonding, that is known to be retained even in solution [24]. Much of the work in quadruply bonded d 4 -d4 systems has involved the group 6 elements (Cr, Mo and W). Almost all these compounds are either neutral or negatively charged. They can be grouped into three classes involving complexes with a) anionic bridging ligands, b) anionic monodentate ligands, or c) compounds with both anionic ligands and neutral donor ligands. A divalent binary trifluoromethylsulfate (triflate) of molybdenum, Mo 2 (SO 3 CF 3 ) 4 , has been obtained by the solvolysis of tetra(acetato) dimolybdenum in trifluoromethanesulfonic  114  (triflic) acid [25]. Characterization by vibrational spectroscopy indicates that the triflate ligands are bridging. This compound has been reported to form complexes with neutral donor ligands such as ethyl acetate (CH3COOC2H5) and acetonitrile (CH 3 CN) [26]. Low valent fluorosulfates are often obtainable by the solvolysis of a suitable precursor in fluorosulfuric acid, HSO 3 F [8]. An analogous synthesis involving Mo2 (CH 3 CO 2) 4 , thus, is a viable route to a low valent binary fluorosulfate of molybdenum. It has been established in previous reports that the oxidation of molybdenum does not produce binary fluorosulfates; instead, the oxide fluorosulfates, MoO(SO 3 F)4 and Mo0 2 (SO 3 F) 2 are produced. The emphasis of this study is on the synthesis of fluoride fluorosulfates of molybdenum by: a) oxidative addition of either bis(fluorosulfuryl) peroxide, S206F 2 , or fluorine fluorosulfate, FOSO2F, to molybdenum pentafluoride, MoF 5 , b) addition of SO 3 F and CH3CN to MoF 5 , and -  c) solvolysis of Mo(CO 2 CH 3 ) 4 in HSO 3 F, analogous to the reaction with HSO 3 CF 3 .  EXPERIMENTAL Reagents The following materials were used in addition to some of the general chemicals described in Chapter 2. Molybdenum powder (99.999%; 200mesh) was obtained commmercially from Alfa Chemicals. Sulfur trioxide (SO 3 ) was obtained from Allied Chemical Corporation, molybdenum hexafluoride (MoF 6 ) from Ozark-Mahonning (now ATO-Chem North America) and acetonitrile (CH 3 CN) from Matheson, Coleman & Bell Manufacturing Chemists; they were all purified by distillation prior to use. Molybdenum pentafluoride (MoF 5 ) was synthesized as reported earlier [27][28] and purified by sublimation. Tetra(acetato) dimolybdenum(II) was synthesized according to literature methods [29].  115  Syntheses Mo0F 3 (SO 3 F): In a typical reaction, 228 mg (1.19 mmol) MoF 5 were allowed to react with  133 mg (0.67 mmol) S 2 06 F2 in a 25 mL Pyrex bulb fitted with a 4.0 mm Kontes stopcock and B10 cone, at room temperature, to give a greenish yellow liquid. The volatiles were removed  in vacuo at room temperature and identified as MoF6 and S 2 05 F2 . The viscous green oily residue was analyzed as MoOF3(SO3F). Analysis found (%): S = 12.08 (calculated (%): S = 11.96). This oil was sublimed quantitatively at 60°C to obtain a white solid having the same chemical composition and exhibiting similar vibrational spectra. The course of this reaction was investigated by  19 F  NMR (vide infra). This product was also formed in the  reaction between MoF 5 and FOSO 2F at ca. 0-10°C, in an analogous manner to the reaction of MoF5 with S20 6 F 2 .  Cs[MoF5 (SQ3al: In a typical reaction, 232 mg (100 mmol) CsSO 3 F and 191 mg (100 mmol)  MoF 5 were weighed out and mixed carefully in a mortar. The intimate mixture was then transferred to a 25 mL Pyrex bulb fitted with a 4.0 mm Kontes stopcock and B10 cone and warmed to 65°C upon which 423 mg of a brown yellow solid were obtained. When left standing at 60°C, the solid gradually turned green. Vibrational spectra show no evidence of either CsSO3F or MoF 5 .  Mo2(SO 3 F)4: 1.5 g (3.5 mmol) Mo 2 (CH 3 C00) 4 were allowed to react with 5 mL HSO 3 F in a 25 mL Pyrex bulb fitted with a B19 cone and a B19-B10 adapter equipped with a 4.0 mm Kontes stopcock. A brownish orange solid was formed as the mixture was warmed to room temperature. After two hours, the mixture was filtered under nitrogen to yield an orange powder that was washed with HSO 3 F and dried in vacuo overnight. This compound analyzed as Mo 2 (SO 3 F) 4 . Analyses found (%): Mo = 32.77, S = 21.91, F = 12.91 (calculated (%): Mo = 32.63, S = 21.80, F = 12.92). The mass spectrum of this compound exhibited a parent ion M + peak at 588 (expected 588).  116  Mo(S03E12(CH3CN)4: In a typical reaction, 141 mg of Mo 2 (SO 3 F) 4 were weighed out in a 25 mL glass bulb fitted with a 4.0 mm Kontes stopcock and B10 cone. Approximately 3 mL of CH 3 CN were added to the brick red powder via transfer in vacuo. Upon warming the mixture to room temperature, a deep blue solution was obtained. After stirring for about 30 minutes at room temperature, the volatiles were pumped off at room temperature for about 6090 minutes, to leave behind 216 mg of a light blue powder. Analysis found (%): C = 20.53, H = 2.73, N = 11.86, S = 14.06, Mo = 21.27 (calculated (%): C = 20.97, H = 2.63, N = 12.22, S = 13.99, Mo = 20.94). The product appears to decompose at 150°C and undergoes another color change at 240°C.  MoF 5 •CH 3 C1V: 191 mg of MoF 5 were reacted with 416 mg of CH 3 CN in an NMR tube fitted with a Young valve. Within minutes of warming up, the yellow powder disappeared giving a clear solution with a slight yellow tint. The volatiles were evacuated after about 1 hour leaving behind a white solid weighing 237 mg. Analysis found (%): C = 10.71, H = 1.33, N = 6.18 (calculated (%): C = 10.35, H = 1.30, N = 6.04).  Reaction of MoF 6 with SO 3 : A 1:1 mixture of MoF 6 and SO 3 was made up at 0°C in an NMR tube fitted with a Young valve. An exothermic reaction took place, at room temperature, giving a mixture with a freezing point of ca. -80°C. The mixture absorbed strongly at the 514.5 nm line of an argon laser. An excess of SO 3 resulted in the formation of a series of poly(sulfuryl) fluorides that were characterized by  19 F  NMR.  DISCUSSION The precursor to some of the compounds synthesized, MoF 5 , is tetrameric in the solid state. The four molybdenum atoms are in a distorted octahedral environment with symmetrical linear Mo-F-Mo bridges [30]. The oxidative addition of S2O6F2 to MoF5 according to Equation 4  117  appeared to be the reaction of choice for the synthesis of MoF 5 SO3 F. This reaction however resulted in the formation of MoF6 , S 2 05 F2 and a viscous oil which analyzed as MoOF 3 (SO3F). The exothermic reaction takes place at room temperature, and does not proceed at temperatures lower than 0°C. The volatile products, MoF 6 and S205F2, were identified by their vibrational spectra and their 19F chemical shifts in the NMR. The progress of the reaction was monitored by  19F  NMR spectroscopy at -10°C using SO 2 FC1 as both a  solvent and an internal reference. The first peak detected was that due to MoF6, and as the reaction proceeded, the intensity of the MoF 6 resonance decreased; peaks attributed to S 2 05 F2 and the fluorine on the SO 3 F group in MoOF3(SO3F) began to appear in the fluorosulfate region (35-50 ppm). The compound MoOF 3 (SO 3 F) was also formed in the reaction between MoF 5 and FOSO 2 F at about 10°C; the other products are MoF 6 and S 2 0 6 F2 . The reaction was monitored by 19 F NMR and experiments indicate that the first species formed are S 2 05 F2 and MoF 6 ; the MoOF 3 (SO 3 F) is the last product to be formed in the reaction. The formation of MoOF 3 (SO 3 F) is common to the oxidative addition reactions of S 2 06 F2 and FOSO 2 F to MoF5 as well as the insertion reaction of SO 3 and MoF 6 . Both the oxidative addition reactions are exothermic and the formation of oxygen has been observed occasionally. It is interesting to note that a clear straightforward addition cannot account for the formation of this compound. There must be some ligand exchange and fluorine migration between molybdenum atoms to genereate a MoF 3 moiety. The yield is approximately 20 25% — the main products being MoF 6 and S 2 05 F2 . It is not possible to formulate a balanced equation for these reactions. Under the experimental conditions,  19 F  NMR studies gave no  clues as to the possible intermediates in the reactions. The viscous oil of the composition MoOF 3 (SO 3 F) could be sublimed quantitatively to yield a white solid of the same composition. The material was studied by vibrational and 19 F  NMR spectroscopy, and may be viewed as an intermediate between the previously  reported Mo0F4 [5] and MoO(SO3F) 4 [10]. The vibrational spectra obtained (Table. 5.1) from the oil and the solid were similar (see Fig. 5-1).  118  Table 5.1. Vibrational data for MoOF 3 (SO3F) (oil)  Infrared  Raman  Assignment  Frequency  Intensity  Frequency  Intensity  1402  w  1393  m  vS02 asym  1175  w  1189  m  vS 02 sym  1139  m  1070  w, sh  1062  m, sh  vS-O  1028  vs  1029  s  vMo=O  865  w  853  m  vS-F  712  s  723  s  vMo-F, vMo-OSO 2 F  675  s  671  m  vMo-F, vMo-OSO 2 F  630  w  621  w  8SO3F  590  w  588  w  SSO3F  560  w  560  m  SSO3F  430  w  450  w  TS 02  326  s  314  m, sh  265  m  202  m  155  w  rSO2F  A  0o U o g 0  j Oil  Solid  I^I^1^1^I  1500^1300^1100^900^700^500 ^  Fig. 5-1. Raman Spectra of MoOF 3 (SO3 F)  ^Wavenumbers (cm')  I  300^100  120  The oxyfluoride derivative, MoOF 4 , is reported to exist in the form of a chain polymer [6][31]. In addition, a metastable structure based on a trimeric unit is also reported to exist for MoOF 4 , with terminal Mo=O bonds and bent Mo-F-Mo bonds resulting in a octahedral coordination about Mo [32]; the bond lengths and angles around the central Mo atom are similar to those found in the chain structure. The vibrational spectra show the presence of both bridging and terminal fluorides. The compound MoOF 4 was also obtained in the present study in a reaction between MoF 5 and HSO 3 F, in addition to HF and S 2 05F2; it was identified by its Raman spectrum [33]. The oxyfluorosulfate species, MoO(SO 3 F)4 , exhibits a distorted octahedral coordination about molybdenum with a terminal Mo=O group and both bidentate bridging and terminal fluorosulfate groups [10]. This compound gives vibrational spectra that are comparable to those obtained for the tungsten analog [11]. The structural similarities between Mo0F3(SO3F) and the two compounds discussed above, MoO(SO 3 F) 4 and MoOF4 are evident from the vibrational data available on the compounds (see Table 5.2), which indicate that all three compounds have terminal Mo=O bonds (vMo=O is between 900 to 1000 cm -1 ). In the case of Mo0F 3 (SO 3 F), the fluorosulfate ligand acts as a bidentate bridging ligand, as seen from the vibrational data, much like the bidentate bridging fluorosulfate group in MoO(SO 3 F) 4 [10]. The three fluorides are all present as terminal ligands similar to the terminal fluorides in MoOF 4 [33]. There are two different kinds of fluorides present which give rise to two overlapping broad signals in the  19 F  NMR at ca. 245 and 233 ppm with an approximately 1:2 intensity ratio; the fluorosulfate resonance is seen as a sharp peak at 41 ppm with a relative intensity corresponding to one F atom (see Table 5.3). Based on the spectroscopic evidence on hand, however, it is not possible to state whether the compound exists in the form of cyclic oligomers or linear chains. It is clear, however, that the molybdenum is in a distorted octahedral environment, in which two of the fluorides attached to molybdenum are equivalent to one another, giving a total of three types of fluorine on the molecule.  121  Table 5.2. Vibrational frequencies of MoOF 4 [33], MoOF3(SO3F), and MoO(SO 3 F) 4 [10] between 1500 and 550 cm -1 MoOF4  M oOF3(SO 3F)  1402  1042  s  w  MoO(SO3F)4  Assignment  1460  m  vSO3  1410  m  vS O3  1243  vs  vSO3  1175  w  1165  w, b  vSO3  1070  w, sh  1045  s, sh  vSO3  1028  vs  1028  s  vMo=0  933  s  vO-SO2F  850  vw, b  vS-F  735  m  vMo-F, vMo-OSO2F  865  w  740  m  721  mw  712  s  vMo-F  668  s  675  s  vMo-F  630  w  636  m  SSO3F  590  w  582  vw  SSO3F  Table 5.3. 19 F NMR data for MoOF3(SO3F)  Multiplicity  Rel. Intensity  Assignment  41  sharp singlet  1  SO3F  233  broad singlet  2  MoF 2  245  broad singlet  1  MoF  Frequency (ppm)  122  A 19 F NMR spectrum and a 95 Mo NMR spectrum of MoF 6 were recorded in the present study. There was a 6-line pattern in the 19 F NMR spectrum due to coupling with the two spin-5/2 isotopes of molybdenum (95 Mo and 97 Mo) in addition to a sharp singlet at  ca. 282 ppm. The value of the coupling constant obtained in the present study (J 19 F-95/97 Mo = 47 Hz) is similar to the value of 44 Hz previously reported by Muetterties and Phillips [34]. Peak integration reveals that 25% of the molybdenum atoms contribute to the 6-line splitting pattern, which indicates that both  95 Mo  (15.72%) and 97 Mo (9.46%) are  coupling to the fluorine atoms. The 6-line splitting pattern (Fig. 5-2a) in the spectrum is remarkably similar to that in a  17 0  19 F  NMR  NMR spectrum of isotopically labelled aqueous  Mo0 4 2- , although in this report [35] the coupling is attributed to 95 Mo only and not to  97 Mo.  A well-resolved 7-line multiplet (1:6:15:20:15:6:1) centred at 562 ppm relative to 1.0 M aqueous Mo0 42- was observed for MoF 6 in the 95 Mo NMR spectrum (Fig. 5-2b). The large quadrupole moments of the two NMR active Mo isotopes usually lead to very broad lines [36], except when very symmetrical environments are encountered about the Mo. The coupling constant J( 95 Mo- 19 F) was determined to be ca. 47 Hz. This value is identical to that obtained from the 19 F NMR spectrum. The solid state reaction between MoF 5 and CsSO 3 F resulted in the formation of Cs[MoF5S03F] according to:  C sS O3F + M oF5  65°C  > Cs[MoF5(SO 3 F)]^-5-  This compound could not be synthesized in the presence of HSO 3 F as a solvent. Other solvents such as CFC1 3 and SO 2FCI were also used but with no success; in all instances where a solvent was used, a viscous material, deep blue in color, was formed. The reaction did, however, take place when an intimate mixture of the two precursors, MoF 5 and CsSO 3 F, was heated to 65°C. An analogous reaction with LiSO 3 F did not proceed, and the vibrational  a)  b)  I i 2 s.,^i 2 $ ^1^R $  1..,,,....1....,...,, ^ ,, .-. , 1-1";^ rprrrrrrr, ,, ;^282 p ^2e1 , 8^281 6 PPM is^It^Ili.^1,  ... .1....,....  r .. ,,, ^  :Ci 8^7112 6^, ;82 •^282  ^Fig. 5-2. a) ' 9 F NMR of MoF6^b) "Mo NMR of MoF6  124  spectrum obtained after mixing the two reagents was a composite of the spectrum of LiSO 3 F and that of MoF 5 . From the vibrational data available on this compound (see Table 5.4) it can be seen that there are five terminal fluorine ligands along with one monodentate terminal fluorosulfate group; the central molybdenum atom is octahedrally coordinated. The diagnostic S-F stretch of CsSO 3 F is conspicuously absent. The tendency of the compound to decompose in solution precludes any NMR studies on this species. The Raman spectrum of the [MoF5(SO3F)] anion -  is comparable to that of its arsenic equivalent, [AsF 5 (SO 3 F)]', obtained by the reaction of C1O2 SO 3 F with AsF 5 [37]. The bands are in good agreement, except in the low frequency  Table 5.4. Vibrational data for Cs[MoF5SO3F] along with selected Raman bands of [AsF5 (S 03F)] [37] -  Raman  C102[A SF5 (S 03F)]  Infrared  Frequency  Mt.  Frequency  Mt.  1395  w  1380  s  1359  s  vS 02 asym  1225  w  1210  s  1205  vs  vSO2 sym  1025  w  1025  s  1034  m  vS-O  813  w, b  800  m  849  m  vS-F  685  s  660  s, b  668  vs  vE-F  580  w  585  s, sh  596  m  bSO3F  562  w  562  m, sh  559  m  8S03F  420  w, b  430  w  451  m, sh  7S02F  320  w  239  w  200  w, b  Frequency  Mt.  Assignment  125  region where ambiguities arise because of the presence of SO 3 F deformation and E-F stretching modes. The insertion reaction between SO 3 and MoF 6 was carried out with the intent of forming derivatives of the type MoF n (SO 3 F) 6 _,I , analogous to the reactions of UF6 and SO3 that have been reported previously. This reaction however resulted in the formation of the polysulfuryl fluorides such as S205F2, S 3 08F2, S 4 0 11 F2 and S50 1 4F2 (see Fig. 5-3). Similar results were obtained from the insertion of SO 3 into SbF5 at 25°C by Gillespie et al. [38].  4E  4E  44  Fig. 5-3. 19 F NMR of polysulfuryl fluorides formed in the reaction between SO3 and MoF 6  The polysulfuryl fluorides were identified by comparison of their chemical shifts in the  19 F  NMR spectrum [38]. Removal of the polysulfuryl fluorides in vacuo left behind a viscous oil that gave a Raman spectrum that was a composite of that of MoOF3(SO 3 F) and SO3. It appears that although some MoOF3(SO3F) was formed in the reaction, it was not entirely free of excess SO 3 . The proposed mechanism is the insertion of SO 3 followed by spontaneous  126  decomposition; the higher polysulfuryl fluorides are presumably formed by the insertion of SO 3 dimers, timers or oligomers. The addition of CH 3 CN to MoF 5 resulted in the formation of MoF 5 CH 3 CN. The vibrational bands in the infrared spectrum of this species are listed in Table 5.5. The average vCN is at 2305 cm -1 , considerably higher than the value of 2267 cm -1 for acetonitrile itself [39]. An explanation for this is the fact that the electron donor pair on nitrogen is weakly antibonding just like the lone pair on carbon monoxide. Hence delocalization of the electrons from nitrogen to the metal is expected to lead to an increase in vCN [40]. The mass spectrum of the compound showed peaks for CH 3 CN and MoF 5 ; the parent peak was not observed. A peak at 149 ppm relative to CFC1 3 was observed in the 19 F NMR, and a broad peak centred at about 6.5 ppm relative to TMS was observed in the 1 H NMR.  Table 5.5. Infrared Spectrum of MoF 5 -CH 3 CN  Frequency  ^  Intensity^Assignment  2950  m  vC-H  2320  s  vC-N*  2290  s  vC-N*  1360  w  vC-C  1015 955 700  s, sh  vMo-F  670  s, br  vMo-F  *The observation of 2 vCN is common for CH 3 CN and many complexes with coordinated  acetonitrile. The cause appears to be Fermi resonance.  127  The reaction between Mo2(O2CCH3)4 and HSO 3 F at room temperature yielded the orange-brown Mo 2 (SO 3 F) 4 according to:  HSO3 F Mo2 (O 2 CCH3)4 + 8 HSO 3 F ^ > 25°C Mo 2 (SO 3 F) 4 + 4 CH 3 CO 2H 2 + + 4 SO 3 F -^-6The infrared spectrum of this compound (Table 5.6) showed bands due to the presence of only a bidentate bridging fluorosulfate group; a Raman spectrum was not obtainable due to poor scattering. The mass spectrum showed a parent peak at m/z = 588; the characteristic pattern of an Mo e fragment is seen down to 198. The reaction is analogous to that carried out with HSO 3 CF 3 to give Mo 2 (SO 3 CF 3 ) 4 , where the presence of a quadruple bond is suggested [25]. The compound Mo 2 (SO3F) 4 is the only known binary fluorosulfate of molybdenum.  Table 5.6. Infrared data for Mo 2 (SO 3 F) 4  Intensity  Frequency  Assignment vSO  1390  s  1335  vs, b  1150  s  vS 02 asym  1055  s  vS02 sym  1030  s  830  s  vS-F  600  vs  vS-OMo  555  s  OSO3F  455  m  TS 03  290  w  rS02  128  Attempts to add S 2 06 F2 across the Mo-Mo moiety met with little success. The addition of CH 3 CN to Mo 2 (SO3 F) 4 resulted in the formation of a deep blue solid of the composition Mo(SO 3 F) 2 (CH 3 CN) 4 . The vibrational spectra of this compound indicate the presence of terminal fluorosulfates only (see Table 5.7). This implies a breakdown of the binuclear structure in Mo 2 (SO 3 F) 4 to give a compound in which there is an octahedral geometry about the central molybdenum where the six sites are occupied by four acetonitrile ligands and two fluorosulfate groups. It is not possible to state whether the fluorosulfate groups are cis or trans to one another based on the evidence at hand. A similar exploratory reaction between CH3CN and Mo 2 (O 2CCH3) 4 was carried out, but after stirring the mixture for one day, there was no weight change in the Mo 2 (O 2 CCH 3 ) 4 upon removal of the volatiles  in vacuo . In an exploratory reaction, ca. 2 mL of HSO3F were condensed onto 241 mg of Mo(CO) 6 in a 25 mL round bottom flask 25 mL glass bulb fitted with a 4.0 mm Kontes stopcock and B10 cone. There was visible bubbling upon warming the mixture, and the color changed from white to yellow-white. The pressure of the volatiles was measured to be 5 mbar at -196°C. Another 2 mL of HSO 3 F were added to the flask and the mixture was left stirring overnight. There was slow bubbling as CO was given off from the mixture overnight. After 24 hours, the pressure of the volatiles was 745 mBar at room temperature. The volatiles were pumped off at -196°C. Upon warming to room temperature there was further outgassing of CO (481 mBar at room temperature), and the volatiles were pumped off at room temperature. A dark green residue (348 mg) was left behind which lost no further weight while being evacuated overnight. The product decomposed at ca. 120°C. The sulfur content was determined to be 24.16%, which corresponds to a composition of Mo(SO 3 F) 3 . The vibrational spectra, however, showed bands at ca. 2133 and 2046 cm -1 in the infrared and at  ca. 2135 and 2076 cm -1 in the Raman indicating that the product was not entirely free of CO. The remainder of the spectrum was poorly resolved. The vCO bands observed in the infrared  129 Table 5.7. Vibrational Spectra of Mo(SO3F)2(CH3CN)4 Infrared'  Raman Frequency  Intensity  Frequency  Intensity  Assignment  2936  w  2938  s  vC-H  2323  m  2323  ms  vC-N  2292  m  2295  s  vC-N  2253  w  2254  w  1367  w  *  1285  m  *  1244  m  vSO3  1170  w  vSO3  1079  m  vSO3  1070  m  *  957  m  *  1067  m  1026  mw  956  w  936  w  *  882  vw  vS-F  722  ms  vMo-O  579  ms  vMo-O  562  m, sh  415  m  415  w  341  m  278  vw  164  vw  'Poorly resolved infrared spectrum makes unambiguous identification difficult; *Bands attributed to CH 3 CN  130  and Raman spectra are at higher frequencies than those observed for other molybdenum carbonyl species [41] and also for Mo(CO)6 itself [39]. There is a coincidence of infrared and Raman bands indicating low symmetry. This reaction bears further investigation. Chromium metal was not oxidizable by S 20 6 F2 at 125°C over a period of four weeks. The addition of a trace amount of CsSO 3 F did not result in oxidation either. The exploratory reaction of uranium metal with S 20 6 F 2 resulted in a species containing a UO 2 moiety as evidenced by infrared spectroscopy, and a reaction between a uranyl chloride and fluorosulfuric acid resulted in complete substitution of the chloride.  CONCLUSIONS Binary fluorosulfates of molybdenum have been a synthetic challenge. The present work is mainly synthetically oriented; mass spectrometry and  19 F  NMR were used  occasionally, and vibrational spectroscopy was used in most instances to characterize the compounds. The work described here has many puzzling results. The reactions of molybdenum pentafluoride with both S 2 0 6 F 2 and FOSO 2F lead to the formation of MoF 6 , MoOF 3 (SO 3 F) and S 2 0 5 F 2 . The reactions do not proceed below 0°C and once initiated, they are extremely exothermic and potentially explosive. It is difficult to postulate a mechanism to explain the formation of MoOF 3 (SO 3 F). What at first appears to be a straightforward viable route to MoF5SO3F, the oxidation of MoF 5 with S 2 06 F2 , is in reality a complex reaction resulting in the formation of, amongst other products, MoF6. The insertion reaction of SO 3 with MoF 6 resulted in the formation of a series of polysulfuryl fluorides in addition to MoOF 3 (SO 3 F). The reaction of MoF5 with HSO3 F, surprisingly, resulted in the formation of MoOF 4 . The only known binary fluorosulfate derivative of molybdenum, Mo 2 (SO 3 F)4 , has been synthesized and characterized in this study. It may be viewed as the fluorosulfate analog of the triflate reported previously. Addition of acetonitrile to this compound resulted in a  131  breakdown of the binuclear species to give Mo(SO 3 F)2(CH3CN)4. It is interesting to note that acetonitrile does not react with Mo 2 (O2CCH 3 )4. It is noteworthy that while a lot of the triflate chemistry of the transition metals is modelled after their fluorosulfate chemistry, in the case of molybdenum some of the fluorosulfate chemistry in the present work has been derived from the previously reported triflate chemistry of molybdenum.  132  REFERENCES  1. E.I. Stiefel, Prog. Inorg. Chem., 22 (1977) 1. 2. 0. Jarchow, F. Schroder and H. Schulz, Z. Anorg. Aug. l Chem., 363 (1968) 58. 3. A.J. Edwards and B.R. Stevenson, J. Chem. Soc. A, (1968) 2503 4. I.R. Beattie, K.M.S. Livingston, G.A.S. Ozin and D.J. Reynolds, J. Chem. Soc. A, (1970) 1210. 5. D.M. Adams and R.G. Churchill, J. Chem. Soc. A, (1968) 2310. 6. D.L. Kepert, "The Early Transition Metals", Academic, New York (1972) p.275. 7. R. DeMarco and J.M. Shreeve, Adv. Inorg. Chem. and Radiochem., 16 (1974) 19. 8. F. Aubke, M.S.R. Cader and F. Mistry, in G. Olah, R.D. Chambers and G.K. Surya Prakash (eds.), "Synthetic Fluorine Chemistry", Wiley, New York (1992) p. 43. 9. J.M. Shreeve and G.H. Cady, J. Am. Chem. Soc., 83 (1961) 4521. 10. K. Lee, Ph.D. Thesis, University of British Columbia, Vancouver, 1980. 11. F.B. Dudley and G.H. Cady, J. Am. Chem. Soc. 85 (1963) 3375. 12. R.E. Noftle and G.H. Cady, J. Inorg. Nucl. Chem., 29 (1967) 969. 13. D.W.J. Cruickshank and B.C. Webster, in G. Nickless (ed.), "Inorganic Sulfur Chemistry", Elsevier, Amsterdam (1968) p. 7. 14. P.A. Yeats, B.L. Poh, B.F.E. Ford, J.R. Sams and F. Aubke, J. Chem. Soc. A, (1970) 2188. 15. H. Imoto and F. Aubke, J. Fluorine Chem., 15 (1980) 59. 16. W.W. Wilson and F. Aubke, J. Fluorine Chem., 13 (1979) 431. 17. R. Dev, W.M. Johnson and G.H. Cady, Inorg. Chem., 11 (1972) 2259. 18. T.A. Stephenson, E. Bannister and G. Wilkinson, J. Chem. Soc., (1964) 2538. 19. J.V. Brencic and F.A. Cotton, Inorg. Chem., 8 (1968) 7. 20. F.A. Cotton and J.G. Norman, J. Coord. Chem., (1972) 161.  133 21. C.L. Angell, F.A. Cotton, B.A. Frenz and T.R. Webb, Chem. Comm., (1973) 399. 22. D. Lawton and R. Mason, J. Am. Chem. Soc., 87 (1965) 921. 23. F.A. Cotton and C.B. Harris, Inorg. Chem., 6 (1967) 924. 24. A.R. Bowen and H. Taube, J. Am. Chem. Soc., 93 (1971) 3287. 25. E.H. Abbott, F. Schoenewolf and T. Backstrom, J. Coord. Chem., (1974) 255. 26. J.M. Mayer and E.H. Abbott, Inorg. Chem., 22 (1983) 2774. 27. R.T. Paine and L.B. Asprey, Inorg. Chem., 13 (1974) 1529. 28. R.T. Paine and L.B. Asprey, Inorg. Synth., 19 (1978) 137. 29. A.B. Brignole and F.A. Cotton. Inorg. Synth., 13 (1972) 81. 30. A.J. Edwards, R.D. Peacock and R.W.H. Small, J. Chem. Soc., (1962) 4486. 31. R.D. Peacock, in J.C. Tallow, R.D. Peacock, H.H. Hyman and M. Stacey (eds.), "Advances in Fluorine Chem., Vol. 7", Butterworths, London (1973) p. 113. 32. A.J. Edwards, G.R. Jones and R.J.C. Sills, Chem. Commun., (1968) 1177 33. D.J. Reynolds, in J.C. Tatlow, R.D. Peacock, H.H. Hyman and M. Stacey (eds.), "Advances in Fluorine Chemistry", Vol. 7., Butterworths, London, (1973) p. 1. 34. E.L. Muetterties and W.D. Phillips, J. Chem. Soc., (1959) 1084. 35. R.R. Vold and R.L. Vold, J. Chem. Phys., 61 (1974) 4360. 36. R.K. Harris and B.E. Mann, "NMR and the Periodic Table", Academic, London (1978). 37. P.A. Yeats and F. Aubke, J. Fluorine Chem., 4 (1974) 243. 38. R.J. Gillespie, J.V. Oubridge and E.A. Robinson, J. Chem. Soc. Proc., (1961) 428. 39. H. Siebert, "Anwedungen der Schwingungsspektroskopie in der Anorganische Chemie", Springer-Verlag, Heidelberg (1966). 40. J. Reedijk and W.L. Groeneveld, Rec. Tray. Chim. Pays-Bas, 86 (1967) 1127. 41. D.M. Adams, "Metal-Ligand and Related Vibrations", Arnold, London (1967).  134  CHAPTER 6. SELECTED REACTIONS OF FLUORINE FLUOROSULFATE *  INTRODUCTION Fluorine fluorosulfate, FOSO 2F, is a fluoroxy derivative or a hypofluorite [1] and, like most of its congeners, is a strong oxidiser. It was first synthesized by Dudley et al. in 1956 [2] by the catalytic (AgF 2) fluorination of sulfur trioxide by fluorine in about 60% yield according to:  220°C F2 + SO3 ^ > FOSO2F AgF 2  Subsequently, the above reaction carried out at high temperatures has been repeatedly reported to give FOSO 2 F in good yield [3]-[5]. The mechanism of the fluorination reaction has been explained by Leung and Aubke [6]. Other methods for the synthesis of FOSO 2 F that have been reported include the thermal fluorination of bis(fluorosulfuryl)peroxide, S 2 0 6 F 2 , by fluorine [7][8], the action of fluorine on NaSO 3 F, Cu(SO 3 F) 2 and Ni(SO 3 F) 2 at 200°C [4], the static reaction of equimolar amounts of fluorine and sulfur trioxide in a stainless steel cylinder at 200°C [9] and the fluorination of sulfamic acid [2]. In addition to the thermal fluorination methods mentioned above, fluorine fluorosulfate has also been obtained by the photochemical reaction of S 2 06 F2 with F2 [10] and OF2 [11], the action of the oxygen fluorides on sulfur trioxide [12]-[14], sulfur dioxide [15] or sulfuric acid [12], and the photochemical fluorination of SO 3 with fluorine [14][16][17]. Very pure fluorine fluorosulfate can be obtained by the thermal decomposition of NF 4 SO 3 F [18] according to:  *The correct name of FOSO2F would be fluoroxysulfuryl fluoride. The name fluorine fluorosulfate first proposed by its discoverers is retained for historical reason. An ionic structure or the presence of F ± or F 6± (SO 3 F) °- is not implied.  135  NF4SO3F  10°C  > NF 3 + FOSO2F^-2-  Fluorine fluorosulfate, FOSO 2 F, is frequently formed as an unwanted and hazardous byproduct in the synthesis of bis(fluorosulfuryl) peroxide [4][5], even though the formation of FOSO2 F requires higher temperatures and/or a higher F2:S03 ratio; however, their different volatilities facilitate their separation by vacuum transfer. The two compounds are interrelated by the reactions: > 2 FOSO2 F^ > 2 S206F2^  F2 + S206F2 FOSO2F + SO3  -3-4-  Fluorine fluorosulfate is a colorless gas with a pungent odor similar to OF 2 . It condenses to give a colorless liquid at -31.3°C, and forms a white solid at -158.5°C. FOSO 2 F does not attack glass, and may be conveniently stored in sealed glass ampoules at liquid nitrogen temperature. Its physical properties, along with those of S 2 0 6 F 2 are listed in Table 6.1. Table 6.1. Some physical properties of FOSO2F and S 2 06 F2 FOSO2 F  S206F2  mp (°C)^-158.5  -55.4°C  by (°C)^-31.3  67.1°C  density (g•mL -1 )^1.770 (69.7°C)  1.726 (25°C)  chemical shift (ppm) * -249, 37  41  Eo-F (1d. mol -1 )^138.072 ± 1.225  98.3 ± 1  vapor pressure curve^log P = 6.56476 - (626.87/T)  log P = 5.49916 (129.25/T) - (259210/T2)  *19 F  NMR, relative to CFC1 3  136  Although the extreme reactivity of FOSO 2F imposed a lower limit of 450 cm -1 (AgC1 plates) when recording the gas phase infrared spectrum [2][19], the use of low temperature techniques has made the complete spectrum available [20]. According to Raman [21] and IR spectra [20] at 298 and 80K respectively, fluorine fluorosulfate has C s symmetry in the liquid and solid state. Of the two possible structural isomers:  the trans form appears to be more stable in the solid state according to a wide-line  19 F  NMR  study [22]. An electron diffraction study on gaseous FOSO 2 F and its Cl and Br analogs is currently being undertaken by Prof. Dr. H. Oberhammer and coworkers (Tubingen, Germany). The kinetics of the reactions of fluorine fluorosulfate have been the subject of a number of studies, in particular the thermal reactions with SO 3 at 90-120°C [23], with SO2 at 60120°C [24], C1 2 at 90-110°C [25], NO 2 at -10 to +10°C [26] and CO at 30-90°C [27][28]. The kinetics of the photochemical reactions of FOSO 2 F with SO2 and SO3 at 2537A from 1525°C have also been studied [29]. In addition to the above reports, an extensive kinetic study of the thermal (236-263°C) decomposition of fluorine fluorosulfate has been reported by Dudley [30]. There are several reviews that briefly cover various aspects of the chemistry of fluorine fluorosulfate. Although some of the reviews are relatively recent, most of the work with FOSO 2 F was carried out in the sixties. Hoffman has briefly reviewed the chemistry of  137  FOSO 2 F [31]. Fokin and co-workers discuss the use of FOSO2F in organic synthesis [32][33]. Some aspects of the inorganic chemistry of FOSO 2F are summarized in reviews by Lustig and Shreeve [34], De Marco and Shreeve [35], Aubke and DesMarteau [36], and Shreeve [37]. At this point, a brief overview of the reaction chemistry of fluorine fluorosulfate may be appropriate. It appears that FOSO 2F may find some use in the one-step synthesis of element fluoride fluorosulfates of the type EF n (OSO 2 F)m , either by oxidative addition to a low valent fluoride or fluorosulfate or by the oxidation of an element, with FOSO 2F acting simultaneously as a fluorinating and fluorosulfonating agent, and there are numerous examples for the concurrent addition of fluorine and a fluorosulfate group. The reactions of fluorine fluorosulfate with olefins and perhalo-olefins have been studied and were found to yield the analogous fluoroalkyl fluorosulfates [38][39], as seen in the following examples:  C 2 F4 + FO SO 2 F  -30°C  C2C14 + FOSO2F  cyclo-0 5 F 8  + FOSO2F  C 4 F 8^FOSO + 2F  25°C 3°C  > CF 3 CF2 OSO2 F  -5-  > CC12FCC12 OS 02F  -6-  >^ cyclo-0 5 F9 OS 02F > CF 3 CF2 CF(OSO 2 F)CF 3  -7-  -8-  Fluorine fluorosulfate reacts with carbon monoxide in the presence of oxygen [27][28] and with fluorosulfonyl isocyanate [40] as shown below:  FOSO 2F + CO FOSO2 F + FSO 2 NCO  02 hv  > FC(0)OSO 2F > FSO2N(OSO2F)C(0)F  -9-10-  The simultaneous addition of fluorine and the fluorosulfate group is also observed in reactions with SO 2 [41], SO3 [5], SF4 [38], and SbF 3 [42] as shown below:  138  SO 2 + FOSO2F SO3 + FOSO2F SF4 + FOSO2F SbF 3 + FOSO2F  195°C 300°C 70°C 25°C  > S 2 05F2  -11-  > S 2 06 F2  -12-  > SF 5 0S02F  -13-  > SbF 40S0 2F  -14-  and the reaction of fluorine fluorosulfate with dinitrogen tetroxide reportedly leads to the formation of nitronium fluorosulfate and nitryl fluoride [43]:  FOSO2F + N 2 0 4  > NO2S03F + FNO2^-15-  The reaction between excess FOSO 2 F and Br 2 results in a slightly volatile yellow liquid of the composition "Br 2 .3FOSO 2 F" [38], similar to the substance obtained by mixing equimolar quantities of BrF 3 and BrOSO 2F:  Br2 + 3 FOSO 2 F  25°C  > "Br2 .3FOSO2 F"^-16-  or by combining BrF 3 and Br(SO 3 F) 3 in various ratios, and a 19 F NMR study of "Br2 .3FOSO 2 F" and such mixtures is reported [38]. However, not all the reactions of fluorine fluorosulfate proceed in such a direct manner. In addition to the straightforward addition reactions listed above, there are also some examples of more complex reactions. The reactions with 1 2 [41] and I(SO 3 F) 3 [38] both result in the formation of trifluorobis(fluorosulfato) iodate(V):  25°C 1 2 + 6 FOSO 2 F ^ > 2 IF3(0S02F)2 + 2 S2 06F2^-17-  ^ ^ 139  25°C I(SO 3 F) 3 + 3 FOSO 2F ^ > IF 3 (0S02F) 2 + S 2 0 6 F2^-18-  The rather broad lines in the 19 F NMR of IF 3 (0S0 2F) 2 indicate that, as in the case of the bromine compound mentioned above, a rapid ligand exchange takes place in the liquid phase. Further examples of the complex reactions of fluorine fluorosulfate are seen with substrates like SOF 2 [41], and Xe [44] which react with FOSO 2 F as follows: 130°C SOF 2 + 2 FOSO 2F ^ > SOF 4 + S 2 0 6 F 2^-19177°C Xe + 2 FOSO2 F ^ > XeF2 + S 2 06 F2^-20-  In all these reactions, fluorine fluorosulfate primarily acts as a fluorinating agent, resulting in the elimination of bis(fluorosulfuryl) peroxide. This shows a distinct difference between FOSO 2F and its Cl and Br analogs, which act only as fluorosulfonating agents towards various halides [35][36] in halogen displacement reactions. Although previous studies provide examples for the diverse behaviour of FOSO 2F, it is unclear from the reports whether, in a given reaction, a straightforward addition is likely to occur or that a more complex reaction will take place, because the studies reported so far provide examples for both straightforward and complex behaviour of FOSO 2F [31][34]-[36]. In spite of the use that FOSO 2F had seen shortly after its discovery, following a cautionary note published in 1968 regarding the unpredictably hazardous and explosive nature of FOSO2 F or an unidentified byproduct [45][46], its use in synthesis has largely stopped. The warning followed a violent explosion at the University of Washington, when a metal cylinder containing a large quantity of FOSO 2F (ca. 1 lb) detonated upon re-opening [47]. We have been interested in the synthesis, structure and reactivity of element fluoride fluorosulfates for some time [42][48][49]. For the most part however, different synthetic  140  approaches have been chosen. The potential importance of FOSO 2 F addition reactions lies in the following areas: a) The synthesis of volatile EF n (SO 3 F) m that have found use in the field of laser isotope separation, b) The generation of unusual cations of the type [EF n] + SO 3 F - , and c) The development of Lewis acids for superacid studies in HSO 3 F [50] and HF. The present study addresses the questions whether FOSO2 F, when carefully purified and used in a cautious manner, is still hazardous, and whether the compound is a useful synthetic reagent, and has a unique reaction chemistry. In order to find answers to the questions raised regarding the use of FOSO2F, and also to provide more complete vibrational and NMR spectra of the principal reaction products, we decided to reinvestigate the previously reported reactions of FOSO2F with SF4 [38], Br 2 [38] and 12 [41], and to expand the scope of this study to the addition of FOSO 2 F to SSF 2 , SeF4 , AsF 3 and MoF 5 . Our previous experience with FOSO 2 F involved its addition to SbF 3 to give SbF 4 SO 3 F in a straightforward reaction [42] as shown above in equation 14. For the element fluoride-fluorosulfates obtained from reactions of FOSO 2F, only limited structural information is available, mainly from 19 F NMR. Only very fragmentary and inconclusive infrared spectra have been reported for most of the EF,(0S02F) In compounds.  EXPERIMENTAL Reagents Selenium dioxide, cesium fluoride, trichlorofluoromethane, iodine, and bromine were obtained from commercial sources and used without further purification. The bromine was stored over phosphorus(V) oxide and KBr prior to use. AsF 3 (Ozark Mahoning, now AtoChem) and SF4 (Matheson) were also obtained commercially and purified by repeated trap-to-trap distillation. Selenium(IV) fluoride was synthesized from Se0 2 and SF 4 as reported previously [51], and MoF 5 was obtained by the reaction of MoF 6 with Mo metal [52]. SSF 2 was prepared by the  141  reaction of sulfur with silver(I) fluoride [53], and SO 2FC1 was prepared by partial fluorination of SO 2 C1 2 . Fluorine fluorosulfate, FOSO 2 F, was obtained (see Chapter 2) as a by-product during the preparation of bis(fluorosulfuryl) peroxide, S206F 2 , although dry ice was used to trap the crude product as recommended [45]. The fluorine fluorosulfate was purified by trapto-trap distillation in vacuo; 19 F NMR was used to test its purity [2] and check for the absence of S206F2. FOSO2F was sealed off in glass ampoules and stored at liquid N2 temperature in a long term storage Dewar (Model 18-XT, Taylor-Wharton).  Reaction Vessels Depending on the amount of reactant used, the exothermicity of the reaction with FOSO 2 F, the anticipated total pressure of the reaction mixture during the reaction, and the desired information on the reaction product, one of the following types of reaction vessels was employed: a)  NMR tubes (5 mm o.d.), fitted with rotationally symmetrical Young valves, for exploratory studies.  b)  Flame sealed glass vials made from standard 6 mm o.d. Pyrex tubing of 25 to 30 cm length for initial exploratory experiments in conjunction with the previously described "ampoule key" [54] to permit opening and re-sealing of the vial in vacuo.  c)  One part Pyrex reactors made from standard or thick wall tubing or from round bottom flasks, and fitted with Kontes Teflon stem valves with an internal volume between 10 and 100 mL as discussed in Chapter 2.  d)  A stainless steel reactor of ca. 150 mL capacity fitted with a Hoke valve for violent reactions (e.g. the reaction of SF 4 and FOSO 2F) or those involving high pressure.  Reaction products were removed from the reactor either by distillation or, in the case of nonvolatile products, by opening the reactor inside the drybox.  142  synthetic Reactions In general, the reactants were combined in the reactor on the vacuum line with the mixture held at liquid N2 temperature. The reactor was then allowed to slowly warm up to room temperature; however, in most instances, the reaction occurred immediately upon warming. The reactions of SSF 2 , SF4 , SeF4 and AsF 3 with FOSO 2F are described in detail. Although the reaction with SF 4 has been reported previously [38], our findings regarding the reactivity of the mixture differ.  a) SSF2 + FOSO 2E: About 1 mmol of SSF 2 was condensed into a reactor of type c), and an equimolar amount of FOSO 2 F was vacuum transferred to this with the reactor kept at -196°C. The reaction between SSF 2 and FOSO 2 F took place explosively while the reactor was held at -196°C. We were unable to increase the scale of this reaction beyond the 20-40 mg range. It was possible to moderate the reaction sufficiently to keep it from occurring below -100°C by adding CFC1 3 to the mixture, but only fragmentation products like SOF 2 , S 2 05 F 2 , SF4 , and a solid identified as S8 were observed.  b) SF 4 + FOSO2E: A stainless steel reactor (type d) was charged with 10 g SF 4 and about 15 g CsF. An equimolar amount of FOSO2F was added in portions of ca. 10 mmol  (ca. 1.2 g) by condensation. On warming, the mixture reacted explosively in some instances and the 0 2 formed was removed in vacuo with the reactor kept at liquid N2 temperature before another portion of FOSO 2 F was added and found to react explosively again. After all the FOSO 2F had been added in this manner, 1.7 g of pure SF5OSO2F (10% yield) were obtained by trap to trap distillation in addition to 02, SF 6 and SO2 F 2 . SF5OSO 2 F was identified by its IR spectrum [55] and vapor pressure [56]. In subsequent attempts, it was noted that the addition of CsF was not necessary. When the reaction was carried out on smaller scales with a 4-fold excess of SO 2F 2 , the formation of oxygen was not observed.  143  c) SeF 4 + FOSO 2F: A mixture of 5 g SeF 4 (ca. 32 mmol) and 1.5 g (ca. 12 mmol) SO2FC1 was placed in a Pyrex reactor (type c) of about 100 mL capacity. To this mixture, 32 mmol of FOSO2 F were added in portions of about 3 mmol by condensation at liquid N2 temperature. Upon melting, the mixture reacted sluggishly and had to be shaken repeatedly. The reaction was monitored by the vapour pressure of FOSO 2F at -100°C [2]. After the final addition, two liquid layers were observed. Trap to trap distillation yielded 1.6 g of pure SeF5OSO 2 F (18% yield) in addition to SeF 6 , solid SeF3OSO2F [57] and some S 2 06 F2 .  d) AsF 3 + FOSO2 F: Typically 150 to 400 mg of AsF 3 were allowed to react in a Pyrex reactor (type c)) with a slight excess (10-20% over the equimolar amount) of FOSO 2 F. Warming the reaction mixture from liquid N2 temperature to room temperature led to a violent reaction which could be moderated by cooling the mixture with liquid N2; occasionally, flashes of light were given off while warming the mixture to room temperature from -196°C. Repeated cooling and warming was necessary before the reaction reached completion and yielded a viscous liquid together with volatile materials identifed as SiF 4 , S 2 0 5 F 2 and S 2 0 6 F 2 , and some 0 2 . The amount of 0 2 formed depended inversely on the cooling applied to the mixture. The reactions of FOSO 2 F with bromine and iodine were performed on a millimolar scale in a similar manner as the reaction with AsF 3 and the results, along with some details, are summarized in Table 6.2 on the following page.  DISCUSSION The four reactions involving SSF 2 , SF 4 , SeF4 or AsF 3 as reactants have been described in detail in the experimental section. They are summarized in Table 6.2 in addition to the reactions with Br 2 and 1 2 , with some emphasis placed on those reaction products that were identified either by IR or  19 F  NMR spectroscopy. While the reaction of FOSO 2 F with SSF 2  explosively leads to decomposition products, the addition to SF4 and SeF4 leads to well-  Table 6.2. Summary of the attempted reactions of fluorine fluorosulfate, FOSO 2 F  No.  Substrate  Reactor Type*  Reaction products, principal product underlined  1.  SSF2  a)orb)  SOF2 , S 2 0 5 P 2 , SF, and S8  -100°C in CFCI,; in pure form explosion at-196°C  Only decomposition and side products are obtained  2.  SF,  d)  SF5OSO2F, S0 2F 2 , SF„ and 0 2  CsF added, S0 2 F 2 formed in situ, see exp. section  Volatile, colorless liquid as previouslydescribed [9][18][19]  3.  SeF,  b) or c)  SeF,OSO,F, SeF, S 2 06F 2 and SeF,OS oy  Moderate reaction at 25 °C; requires shaking of reaction mixture  Volatile, colorless liquid  4.  A sF,  a)b)orc)  AsF 6(S0,95 , SiF4 S 2 0 5 F 2 , S 2 06F2 and 0 2  Explosive, vigorous; requires slow careful addition ofFOSO 2F  Viscous liquid oflow volatility n-3.5; similar to AsF,(SO,F),  5.  MoF,  b)  Mo0F,(0SO2F), M 0 F6 and S205F2  Slow reaction; similar to the MoF 5 -S 2 06 F2 reaction (chap. 4)  Viscous liquid may be sublimed in vacuo to givelow melting solid  6.  SbF,**  c)  SbF,OSO,F  Slow reaction at25°C over8 days  Viscous, colorless liquid [6]  7.  Br,  c)  BrF n(SO,F),, SiF 4 , 0 2 S 2 0,F, and S206F2  Vigorous, occasionally explosive reaction at room temerature  Viscous, pale yellow liquid, as described before [9], n-1  8.  12  IFXSO,F)„,, SiF4 , 0 2 S 2 05 F 2 and S206F2  Vigorous; flames are visible; goes through yellow intermediate  Viscous, clear, colorless liquid distillable [10], n-3.3  a)or c)  *See Experimental section for explanation **Details from W.W. Wilson, Ph.D. Thesis, UBC, 1975  Comments and observations on the reaction  Description ofprincipal products  145  defined compounds of the type EF5SO 3 F (E = S, Se), and the reactions with AsF 3 and I 2 yield non-stoichiometric compounds of the type EF n (OSO 2F) 5 _n (E = As, I; n between 3 and 4). In this section we will discuss the compounds and their vibrational and NMR spectra. In order to elucidate the course of these rather complex reactions which yield a range of different products, we will turn our attention to the byproducts. It appears that in the reactions carried out, the oxidative addition of FOSO 2 F was always accompanied by side or secondary reactions. All the reactions studied proceeded vigorously, even explosively in some instances. Our observations suggest the following order of reactivity towards FOSO2F under the conditions employed here: SSF 2 > > SF 4 > Br2 > 12 > AsF3 - MoF5 - SeF 4 Attempts to moderate the reactions by adding a third component as a dilutant to the starting mixture have had limited success because the additives used, CFC1 3 and also SO 2FC1, are degraded under the reaction conditions. The use of SO2 F2 as a dilutant results in high internal pressure, and the reactions may only be carried out safely in a metal reactor.  The stoichiometric fluoride- fluorosulfates Although the mixture of SSF 2 and FOSO2F reacts explosively at -196°C, the reaction can be moderated to some extent by the addition of CFC1 3 . The reaction of SSF 2 with FOSO 2 F at -100°C in CFC1 3 may be formulated as:  2 SSF 2 + 4 FOSO2 F  > 2 S 2 05 F 2 + 2 SOF 2 + SF 4 + 1/8 S8  -21-  but a detailed reaction mechanism is, however, lacking. The materials obtained from the reactions with SF 4 and SeF 4 , albeit in rather low yield, are volatile, well-defined, stoichiometric compounds of the composition EF5OSO 2F. Both compounds were easily purified, and their compositions were established by vapor density and spectroscopic measurements. There was a remarkable difference in reactivity  146  between SF 4 and SeF 4 towards FOSO2 F, which was reflected in the yield of the products. In the case of SF5OSO 2F, the yield was around 10% (lower than the 25% reported previously, where a temperature of 70°C and a reaction time of 3 hours were required [38]). The yield of SeF5 OSO2 F was considerably higher (18%) and decomposition products such as 02 , SO2 F2 or S 2 0 6 F 2 were not observed. The formation of a white solid residue, identified by its IR spectrum as SeF 3 OSO2 F [57], in the rather slow reaction between SeF 4 and FOSO 2F is puzzling. It appears that the oxidation of selenium to the +6 oxidation state may not occur readily. While it is possible to formulate the overall reaction as:  3 SeF4 + 2 FOSO2 F ^ > SeF 5 OSO2 F + SeF 3 OSO2 F + SeF 6^-22-  which accounts for all the observed products, the above reaction does not clearly explain how they form. It is also possible to explain the formation of SeF3OSO2F by a ligand exchange reaction of the type:  SeF5O SO2F + SeF4 ^ > SeF6 + SeF3OSO2F^-23-  It appears that alternate published routes to SF 5 OSO 2F such as the photolytic reaction of sulfur dioxide with bis(pentafluorosulfur) peroxide [55] and the direct fluorination of sulfur in the presence of oxygen or oxygen fluorides [56] may be more efficient and less hazardous. This conclusion, however, cannot be extended to the synthesis of SeF 5 OSO2 F described here. The previously reported route, the photolysis of a mixture of (SeF 5 ) 2 02 and S206F 2 [58], does not offer any advantage on account of the thermal and photolytic instability of (SeF5)20 2 . The complete vibrational spectra of the monomeric, well-defined compounds of the composition EF 5 OSO2F (E = S or Se) are compared to the previously reported IR and Raman spectra of TeF5OSO2F [59]. The data for all three compounds are listed in Table 6.3 together with estimated intensities, the depolarization ratios for the more intense Raman bands and an  ^  Table 6.3. Vibrational Spectra of the Pentafluorofluorosulfates of Sulfur, Selenium and Tellurium in the Frequency Range of 1500-100cm -1 SF,OSO,F Infrared(g)  SeF,OSO,F  Raman (I)  Infrared(g)  r[cml  int.  6,1, [m►► l ]  int.  1493  s  1490  03  -P  1255  s  1252  5  0.02p  888  vs  885  0 .2  <0.1P  s 830 -830 ^vs 947 907 m 910 s 848 w,sh 776 m 730 733  0.5  673 645 620 604 575 553 503 475 418  m w,sh w,sh m ms w,sh mw w w  p  vfrnf i l  in  1492 1453  vs w  872  vs  TeF,OSO,F*  Raman(1)  Infrared(g)  Approximate band description  Raman(1)  Av [cm'] 1486  Mt. 0.8  p 0.5p  v [mil 1490  Mt. vs  Av [mul 1476  int. m  p  1249  4  0.05p  1250 1160  vs vw  1242 1150  s vw  P  880  1  0.5p  895  vs  893  vw  vs <0.1p 851 845 F vs 772 770 s <0.1p 746 0.1 744  1.5  0.1p  850  vs  844  m  p  05 0.5  P p  750 730  vs vs  728  m  p  10 0.01 6.4  "OP dp P  681  vs  P  EF,0^v w,,,^AI  641  s  P  EF,0^v.,..^A, uncertain possibly E SF50  SO,^v. SO,^vv,,, .0 v_ - F^ o^ v s ^""' EF5O^v,.^B s  1  EF,o^v.^E  10  Op  675  w  675 650 632  0.5 03 0.5  <0.1p -Id <03p  647  m  580  0.1  P  593  m  590  0.1  P  578  m  579  s  555  0.1  P  541  wsh  540  0.1  P  545  w  532  w  SO,^8,  483 420 304 286 266  0.2 ~1 dp 0.7 0.3p 0.3P 2.2 0.5 -I dp I S^0.2p  1.0 0.1 1.6 4  0.3p 0.5p 0.5p 0.2p  448  m  w  460 424 335 291  EF5O^8. SOS^rock.  s s s  160  0.2^0.5p  112 2 0.2  0.5p 0.2p  315 277 263 175  m  676 660 645  641  s  ,  460 333  *ref. 24, IR spectra recorded between NaC1 windows  230 217 140  P  SO,^8.  unassigned deformation and torsionvibrations  148  approximate band description. The infrared spectra of SF 5 OSO 2F and SeF5OSO 2F are shown in Fig. 6-1, and the Raman spectrum of SF 5 SO3 F is displayed in Fig. 6-2 to show the polarization ratios of the bands. A previous report on the IR spectrum of gaseous SF5OSO 2 F [55] lists a limited number of bands in the region from 1500-550 cm -1 , with band positions in good agreement with our findings. It is apparent from the listings in Table 6.3 that all three EF 5 OSO2 F (E = S, Se and Te) compounds have a monodentate fluorosulfate group covalently bonded to an EF5-moiety. As expected, the SO 3 F-vibrations show little variation with E, while mass differences between S, Se and Te cause a gradual shift of EF5-vibrations to lower frequencies. A local C4 v symmetry is suggested for the EF 5 - moiety, and four stretching vibrations (E, 2A, B) are expected. The assignments previously reported for TeF5OSO2F [59] are retained here. Very intense polarized Raman bands, assignable to the symmetric stretching motion of the group, are observed at 733 cm -1 for SF5 OSO2F, 676 cm -1 for SeF5OSO2F and 581 cm -1 for the Te analogue, in agreement with local C4 v symmetry of the EF5O- moiety. Additionally, some weak and probably depolarized Raman bands are attributed to this group. In contrast, all the SO 3 F-vibrations are polarized for the three compounds. While the symmetric EF5O stretching vibrations are observed for all three EF5OSO 2F compounds about 20-30 cm -1 lower than the A ig mode in the corresponding hexafluorides [60], the frequency of the highest IR active EF5O-stretch is in about the same position as v 3 (F lu) for the EF 6 compounds. This band is not observed in the Raman spectra of TeF5OSO 2 F and SF5OSO2F. Additional instances of the mutual exclusion of IR and Raman bands are found in Table 6.3. The near octahedral symmetry observed in the EF5Omoiety of the monofluorosulfates also explains why the EO-vibrations are not clearly identifiable, but are instead mixed in with the EF 5 -vibration. There are two general areas of ambiguity regarding the attempted group frequency assignment in the stretching region: For the monodentate -OSO2F group, a clear distinction between vS-F and vS-O is not possible; both are observed in the narrow region of 890-  149  SeF5SO3F  1030  1454  671  956 896  646  595  529  747  '1  1492^1250  872 1  850  426  771  SF 5 SO 3 F 776 v673V 731 668 605  1256 1492 828  575  888 948  1600.0  1431.2 1262.5^1093.7^925.00 756.25^587.50 418.75 250.00  Wavenumbers (cm -1  )  Fig. 6-1. Gas phase infrared Spectra of SF5SO3F and SeF5SO3F  A  1 733 1251  3041 266 I 832  Wavenumber (cm -1 ) Fig. 6-2. Raman Spectrum of liquid SF 5 SO 3F  420  151  840 cm -1 , and the vibrations are expected to couple. Hence, an assignment as asymmetric and symmetric stretches of an 0-S-F moiety has been used. In addition, vSF 5 vibrations are expected to occur in this region as well, resulting in both partial band overlap as seen in the 830-840 cm -1 region of the Raman spectrum, and in vibrational coupling and extensive band proliferation in the S-F and S-0 stretching region (600-950 cm -1 ). These complexities are absent in the spectra of SeF5OSO2F and TeF5OSO 2F [59]. A final comment concerns the SO2F-stretching vibrations. There is generally very good agreement between the IR band positions, obtained from gaseous molecules and the Raman bands found in the spectra of the liquid compounds. For TeF 5 OSO 2F, small discrepancies are noticeable, with Raman bands found at slightly lower frequencies than their infrared counterparts. Very weak intermolecular association in the liquid phase via fluorosulfate groups is the most plausible cause for the gas to liquid phase shift. The high frequencies of the vSO 2 vibrations are typical for covalent, monodentate fluorosulfate groups (e.g. FOSO 2 F or S 2 06 F 2 [21]), where frequency lowering of the E-OSO2 F band because of intermolecular association or increasing ionicity of the E-OSO 2 F bond can be ruled out. A very slight shift to lower frequencies is noted when comparing both the SO2 stretching vibrations in the EF 5 OSO2 F compounds ( S < Se < Te) and suggests increasing polarity of the E-OSO 2 F bond. The 19 F NMR spectra of SF5OSO2F and SeF5OSO 2F have been reported previously [55][581161], and are in good agreement with this work. One expects to find a strong similarity between the 19 F NMR spectra of SF5OSO2F and its selenium analogue, both of which exhibit AB 4 X spectra. In addition, satellite resonances due to coupling of 19 F to 77 Se are observed in the 19 F NMR spectrum of SeF5OSO 2 F. The 77 Se NMR spectrum shows a six line pattern of relative intensity ratio 1:5:10:10:5:1 which implies that J( 77 Se-19 Fax) .?--• J( 77 Se- 19 Feq ). Small splitting of the individual lines is indicative of second order effects; there is a difference of about 20 Hz in the coupling constants.  152  The non-stoichiometric fluoride- fluorosulfates The viscous materials of low volatility produced when Br 2 , 12 or AsF 3 , are reacted with FOSO 2 F are non-stoichiometric. Their exact composition depends on the reaction conditions and the amounts of reactants used. This observation contradicts earlier reports on the reactions of Br2 and 12 with FOSO 2 F, where compositions like Br 2 .3FOSO 2F [38] and IF 3 (0S02F)2 [41] were claimed. For the reaction with iodine, compositions ranged from IF 3.3 (0S0 2 F) 1.7 to IF 3 (SO3 F) 2 , and with bromine, the approximate composition BrF(OSO 2 F) 2 was obtained. Here, and in the case of the reaction product formed from AsF 3 and FOSO2F (approximate composition AsF 3. 5(0S0 2F) 1. 5), the SO 3 F content is higher than anticipated for either addition of FOSO 2 F or fluorination. This implies that the S 2 06 F2 formed as a byproduct is also involved in the oxidation process, while some of the fluorine is converted to SiF 4 during the highly exothermic reaction which induces fluorine attack on the glass wall. Attempts to detect either AsF5 or any bromine(V) fluoro species by either 19 F  NMR or vibrational spectroscopy were unsuccessful; however, both compounds, if  formed, would have reacted with glass under the reaction conditions. The reactivities of AsF3 and SbF 3 [42] towards FOSO2F exhibit a similar trend as compared to the group 16 elements, i.e. AsF3 > SbF 3 . Some of the flashes of light observed during the reaction of AsF 3 with FOSO2F may be due to warming AsF 3 from liquid N2 temperature which reportedly causes greenish-yellow luminescence and acoustic emission [62]. A different bonding situation for the fluorosulfate group is encountered for IF„(0S0 2 F) 5 _„ (n 3) and more so for BrF n (OSO2F) 3 _n (n = 1), as compared to the EF5 SO3 F discussed earlier. The Raman shifts are listed in Table 6.4. Due to the high reactivity of the materials, infrared spectra could not be obtained. A relatively simple Raman spectrum is obtained for IF n (OSO 2F) 5 .. n . The vSO2 at 1424 and 1212 cm -1 are at considerably lower frequencies than those observed for EF5OSO 2 F, probably caused by weak intermolecular association. A slightly polar monodentate fluorosulfate group is suggested, consistent with vS-O and vS-F at 940 and 857 cm -1 . A very similar Raman spectrum is  153 Table 6.4. Raman Spectra of Br(SO3F)3, BrFn(SO3F)3- n and IFn(SO3F)5-n  Br(SO,F),(s)* Av(cm-1 ) 1490 1467 1372 1356 1241 1230 1168 1122 1010 859 827 801 721 645 612 583 563 551 540 455 430 408 384  BrFn(SO3F),(1)  1F.(SO3F),..(1)  HSO,F**  Int.  Av(cm 1 )  Int.  mw w m mw, sh s mw, sh mw m m  1480 1460  mw w, sh  1424  m,br  1443  m  1360 1240  m,br s  1212  s  1205  s  1140 1065 1030 920 850 830  w,br w m vw m w,sh  940 857  vw ms  961 851  w m  750 650 610 580 —550  s,br vs vs ms m, sh  706 652 612 585 555  vs vs vs ms m  560 552  m m  455  s  460 435  408  m ms  mw  490 405  m w  395  m  393  m  m vw m ms vs,br vs ms m vw w s w mw mw  303  vs  276 225 206  vs ms ms  Av(cm I )  hit.  305^} 300 260 220  s  280  s  ms m  257  vs  150  w,sh  160  w  *ref [25]. **ref [11], bands due to the OH group like OH-stretching (2940 cm -1 ), OH-bending (1179 cm") and OH-torsion (686 cm') are omitted.  Av(cm- ')  Int.  154  obtained from the fluorosulfate group in HSO 3 F [21] (included for comparison in Table 6.4). The good correspondence in the band positions also extends into the region of the deformation modes. Very prominent and intense Raman bands are observed at 706, 652 and 612 cm -1 ; these are not assignable to the fluorosulfate group. A comparison to the Raman spectrum of IF5 [60], which displays strong bands at 710, 693 and 605 cm -1 , might suggest an assignment as I-F stretching vibrations. However this assignment for the bands at 652 and 612 cm -1 is ambiguous, because for I(SO 3 F) 3 and the salts M[I(SO 3 F) 4 ] (M = Na or K) very strong Raman bands are observed between 640 and 612 cm -1 [63], and may involve iodine-oxygen stretching motions. A very similar situation is encountered for BrF n (OSO 2 F) 3 ,  (n = 1) where very intense  Raman bands in the same region (650 and 610 cm -1 ) are also found for Br(SO3F)3 [63] (included for comparison in Table 6.4). A single Br-F stretch is expected for a material having the composition BrF(OSO 2 F) 2 , and an assignment of vBr-F is suggested for the broad band found at 750 cm -1 . The similarity between the Raman spectra of BrF(OSO2 F) 2 and Br(SO 3 F) 3 extends over the entire spectral range. It is not unexpected to see a greater degree of band splitting in the case of Br(SO 3 F) 3 , which is a crystalline solid. It is hence reasonable to suggest oligomeric structures which imply the presence of bridging, bidentate and terminal, monodentate fluorosulfate groups for both BrF(OSO 2F) 2 and Br(SO 3 F) 3 . Formulation as BrF(OSO 2 F) 2 is preferred here because no single, preferred bond type is likely here, unlike it had been in EF 5 OSO 2 F (E = S, Se or Te). It is also noted that this spectral similarity extends to the Raman spectra of I(SO 3 F) 3 [63] and Au(SO 3 F) 3 [64] where, in all instances, a nearly square-planar coordination environment of the central atom is suggested. This assumption is confirmed by an X-ray diffraction study, in the case of [Au(SO3F)3] 2 [65]. Intense Raman bands at 650, 450 and 280 cm",1 are observed for BrF(OSO 2F) 2 as well as for the tris(fluorosulfates) of bromine, iodine  155  and gold. These vibrations involve, in part, E04 -skeletal vibrations (E = Br, I [63] or Au [64]) or, as in this case, BrO 3 F-skeletal vibrations. Similar to the situation above, where a nearly octahedral EF5O- environment exists in EF5OSO 2 F (E = S, Se or Te), the substitution of a single F by 0 (or in this case 0 by F) that involves atoms of similar mass and bonding characteristics does not appreciably alter the vibrational spectrum of the still fairly symmetrical coordination environment around the central atom. Little needs to be said at this point about the vibrational spectra of the remaining fluoride-fluorosulfate species discussed in the present study. The Raman and IR spectra of SbF 4 OSO2F have been reported and discussed in terms of a bidentate, presumably bridging fluorosulfate group [42], and those of OMoF3(OSO 2 F) have been included in an extensive study of the molybdenum oxo-fluoro- fluorosulfato derivative in Chapter 4. The Raman spectrum of AsF,,,(0S0 2 F)5_n (n = 3.5), is largely identical to the one observed previously for AsF 3 (OSO 2 F) 2 [48]; a new band is observed in the As-F and As-0 stretching region at 684 cm -1 in addition to previously reported bands, which are very slightly shifted relative to those reported for AsF 3 (OSO 2 F) 2 and which are quoted in brackets 754 s (750 s) 669 vs (670 vs), 645 s, sh (652 vs, sh) and 610 vw (610 vw). It is tentatively suggested that the product obtained from the reaction of AsF 3 and FOSO 2F is a mixture of AsF 3 (OSO 2 F) 2 , formed when S 2 06 F2 reacts with AsF3 [42], and AsF 4 OSO 2 F, consistent with its non-stoichiometric composition. Much like the vibrational spectra of the five compounds discussed here are different from those of the EF 5 SO 3 F, the 19 F NMR spectra are also different for the nonstoichiometric, low boiling or non-volatile products of the reactions of FOSO 2 F with AsF 3 , 12 and Br2. The spectra differ slightly, depending on the experimental conditions, from one preparation to the next in chemical shifts, number of peaks, peak shape and peak integration ratio. The pertinent data are summarized in Table 6.5. For AsF n (OSO 2 F) 5 _„, two broad resonances were observed for one sample. One of the resonances had an almost identical chemical shift of -37.8 ppm when compared to the reported  ^ 156  Table 6.5. Characterization of the fluoride-fluorosulfates of sulfur, selenium, arsenic, bromine and iodine by NMR. Compound SF,SO,F  SeF,SO,F  5(1V111)*^Int.^Assignment 72.8^d^4^Fe, 56.8^q^1^F. 46.3^s^1^F(sc,3F)  4F.-F.4)=155Hz  74.9^d^4^F., 58.3^q^1^F. 46.9^s^1^Foo3,..,  J(F-F,,,)= 217 Hz  8"Se=-675ppmrel.Se0,inf1 2 0/D 20 J(n Se-F.,) — J("Se-F.,)— 1430 Hz AsFe(SO3F)5_o  BrFe(SO3F)3.e** n-0.84to1.13 IFe(SO,F),„ n=2.89 n=2.65 n=3.32-3.17 n=1.63  (i) -25.6^1^As-F -37.8^2.5^As-F 43.8^ F(s,33,) (ii) -41.4^broad^As-F 42.5^ F(so3F) -46.2^to-48.8^Br-F 45.8^to47.9^F(so3F) (i) -14.3^and-18.3^I-F 45.7^ F(so3F) (ii) -15.4^and-19.6^I-F 47.3^ F003,, (iii) -18.2^and-2.09^I-F 48.7^ F(s03,, (iv) -29.6^ 51.1^  *rel to CFC1 3 external standard, recorded unlocked * *values from three different samples  I-F F(s03,)  157  value of -37.2 ppm for AsF 3 (OSO 2 F) 2 , [42], which suggests two different arsenic-fluorine compounds in the non-stoichiometric mixtures, one of them being AsF3(OSO 2 F) 2 . Like in AsF 3 (OSO 2 F)2, [42] only a single resonance in the fluorine-sulfur region was observed with (5 = 43.8 ppm, compared to the value of 45.6 ppm reported for AsF 3 (OSO 2F) 2 [42]. In another preparation, only a single broad As-F resonance at -41.4 ppm was observed. Single resonances are found for both the BrF and the SO 3 F groups in BrF n (OSO2 F) 3 _n • This is in agreement with observations by Gilbreath and Cady who used mixtures of BrF 3 and Br(SO 3 F) 3 to study the entire composition range and reached similar conclusions regarding the non-stochiometric nature of these mixtures [38]. There is only a slight discrepancy that concerns the chemical shift range for the single fluorine on sulfur resonance; our observed values between 45.8 and 47.9 ppm are higher than the range of 39-41 ppm reported [38]. The 19F NMR spectra of IF n (OSO 2 F)5_ n show the greatest variation, depending on the course of the reaction. Either one or two broad resonances are found in the I-F region and peak positions vary with composition. Previous reports [38][41][66] had differed on the number of resonances in the I-F region and on the peak integration ratios. It seems now that the previous studies using different synthetic routes, may have produced materials of varying composition rather than well defined stochiometric compounds as claimed. Considering the different shapes of the resonances, and the broad peaks found in particular for the IF,- and AsF n - fluorosulfate species, peak integration is ineffectual, and compositions determined in this manner are very approximate. Frequently repeated measurements on the same compound, as e.g. the IF n (OSO2 F) 5 _„ (Table 6.5, iii), give different values for  n.  Unresolved 19 F- 19 F splitting or coupling to the quadrupolar nuclei I (I = 5/2) and As (I = 3/2), both of 100% natural abundance is a likely cause for the broad resonances. It is unclear whether the two broad resonances observed frequently for both materials implies the presence of different species (e.g. AsF3- or AsF 4 -fluorosulfates), as suggested by the Raman  158  spectrum of AsF n (OSO2F)5„, or are due to different fluorines in the same molecule (e.g. F in an equatorial or axial site of an approximately square pyramidal iodine(V) in IF n (OSO 2 F) 5 _n)• A common feature of all three non-stoichiometric fluoride-fluorosulfates is the observation of a single  19F  resonance for the fluorosulfate group. For BrF n (OSO2 F) 3 _n rapid  exchange has been suggested to account for this feature, and various ionic fragments formed by dissociation of an oligomeric substance have been proposed [38]. Similar arguments can be made for IF n (OSO 2F) 5 , and AsF n (OSO 2F)5_n as well. In summary, the vibrational and more so the  19 F  NMR spectra support the contention  that the reactions of Br2, 12 and AsF 3 with FOSO 2F produce non-stochiometric oligomeric mixtures of varying composition, depending on the mole ratios of reactants used and the reaction conditions.  The decomposition of fluorine fluorosulfate The formation of the byproducts in all the reactions can be explained by examining the decomposition modes of FOSO 2F and the ensuing products (Fig. 6-3). The in situ generation of SO 2 F 2 , together with that of 0 2 , during the addition of FOSO 2 F to SF 4 as described in the experimental section, may be explained by the decomposition of fluorine fluorosulfate according to:  FOSO2F ^ > 1/2 02 + SO2F2^  -24-  The kinetics of this decomposition in the temperature range of 509-536K have been studied by Dudley [30], who found that the reaction is catalyzed by the addition of fluorine. In the absence of fluorine, an alternative decomposition reportedly occurs [30] via intermediate SO 3 F . radicals according to:  FOSO 2F ^ > 1/2 S 2 0 6 F 2 + 1/2 F2^  -25-  159 02 +50,F,  1 F-0--SO,F  I, F + *OSO,F .  S206F2  F.B. Dudley, J. Chem. Soc.,(1963)1406.  1  S20 5F 2 + 0 2  Fig. 6-3. Decomposition scheme of FOSO 2 F [30]  This reaction may explain the formation of bis(fluorosulfuryl) peroxide during the reactions of SeF 4 , Br2, 12 and AsF 3 with a slight excess of FOSO 2 F, if it is assumed that the F2 formed will react further immediately. In three of the above reactions, large quantities of  SiF4 were observed. Some SeF 6 was also produced in the reaction with SeF 4 (which proceeds rather slowly), presumably by fluorination of unreacted SeF 4 . It can be argued that FOSO 2F acts primarily as a fluorinating agent in these cases; however, when Br 2 is used as a reactant, any 5 2 06 F 2 formed is capable of oxidizing it as well, resulting in the formation of BrF,i (OSO 2 F) 3 ,  (n --= 1).  The simultaneous formation of S 2 06 F2 , S 2 05 F2 and 02 , together with SiF 4 , suggests another probable side reaction which has been previously observed in quartz cells above 120°C [66] — the decomposition of SO 3 F radicals to give S 2 05F 2 and 02 according to: .  S03F'  > 1/2 S 2 05 F 2 + 1/2 0 2  -26-  160  The reactions studied here are sufficiently exothermic to bring about temperatures high enough to allow the secondary reactions suggested above to occur.  CONCLUSIONS In summary, fluorine fluorosulfate (FOSO 2 F) is extremely reactive, leading to very exothermic and frequently explosive reactions with various substrates. The simultaneous oxidative addition of fluorine and fluorosulfate is almost always highly exothermic, and results in side reactions and thermal decomposition of the excess FOSO 2 F, thereby markedly reducing the yield of the expected products. It is somewhat surprising to us that in earlier uses of FOSO 2F in synthetic chemistry, the explosive nature of this compound, the complex course of its reactions, and the non-stoichiometric composition of some of its reaction products had not been fully recognized [31][34]-[36]. The reactions involving the addition of FOSO2 F to SeF 4 and, as reported previously, to SbF 3 [42] are useful synthetic routes, and while it is possible to obtain some useful and interesting information on the element fluoride-fluorosulfates, the use of fluorine fluorosulfate in their synthesis is questionable and best avoided. When available, alternative routes to these compounds involving bis(fluorosulfuryl)peroxide, S 2 0 6 F 2 , are therefore preferred. As demonstrated recently, when ligand exchange is observed, it can be exploited very well as shown by the facile one-step synthesis of fluoride fluorosulfates of niobium and tantalum having the general formula MF,i (OSO2F)5_, [68]. In most instances, the risks in preparing, storing and handling, and using FOSO 2F as synthetic reagent far outweigh its rather limited usefulness, and the words of warning by Cady [45] have not lost their validity. The results of the present study strongly support these comments regarding the hazardous nature of fluorine fluorosulfate, FOSO 2 F, and this compound must be included in the same class of highly dangerous and unpredictable chemicals along with compounds such as FONO 2 and FOC10 3 [34].  161  REFERENCES  1. G.H. Cady, XVIlth IUPAC Congr. Vol. 1, lnorg. Chem., Butterworth, London (1960) p. 205. 2. F.B. Dudley, G.H. Cady and D.F. Eggers, Jr., J. Am. Chem. Soc., 78 (1956) 290. 3. G.H. Cady and F.B. Dudley, US Pat. 2876075 (1959), Chem. 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Chem., 47 (1969) 4153. 58. W.L. Reichert and G.H. Cady, lnorg. Chem., 12 (1973) 769. 59. H. Burger, Z. Anorg. Allg. Chem., 360 (1968) 97. 60. K. Nakamoto, "Infrared Spectra of Inorganic Compounds", 2nd edn., John Wiley and Sons, New York (1970), and references therein. 61. C.I. Merrill, S.M. Williamson, G.H. Cady and D.F. Eggers Jr., lnorg. Chem., I (1962) 215. 62. H. Russell Jr., R.E. Rundle and D.M. Yost, J. Am. Chem. Soc., 63 (1941) 2825. 63. H.A. Carter, S.P.L. Jones and F. Aubke, Inorg. Chem., 9 (1970) 2485. 64. a) K.C. Lee and F. Aubke, Inorg. Chem., 18 (1979) 389. b) K.C. Lee and F. Aubke, lnorg. Chem., 19 (1980) 119. 65. H. Willner, S.J. Rettig, J. Trotter and F. Aubke, Can. J. Chem., 69 (1991) 391. 66. F. Aubke and G.H. Cady, lnorg. Chem., 4 (1965) 269.  164  67. F.B. Dudley and G.H. Cady, J. Am. Chem. Soc., 85 (1963) 3375. 68. D.L. Zhang and F. Aubke, J. Fluorine Chem., 58 (1992) 81.  165  CHAPTER 7. SUMMARY AND CONCLUSIONS  GENERAL COMMENTS The systems studied in this thesis and the reagents used in these studies encompass a very wide range. New anions of the early transition metals (Ti, Zr, Hf and Mo), where the metals were primarily in their highest oxidation state, were isolated by using Cs + to stabilize them, as discussed in Chapters 4 and 5. Selected organotin(IV) compounds were characterized by Xray crystallography in Chapter 3. The principal reagents and solvents used (HSO 3 F, S 2 06 F2 and FOSO 2F) belong to the non-metallic groups 15, 16 and 17, as do some of the compounds studied in Chapter 6 (MF 4 (M = S or Se), AsF 3 , Br2 and 1 2). In addition, unusual oxidation states of electron-rich gold compounds are discussed in Appendix A. While the casual reader may get the distinct impression that we are taking a random walk through the periodic table, this chapter will demonstrate that our travails were not random by any means, but were indeed, purposeful. There are two complementary themes which connect and pervade all these, at first seemingly unrelated, topics. The first theme is that of superacidity, with emphasis placed on conjugate superacids in fluorosulfuric acid, HSO3F. One facet of this theme is the development of new Lewis acids in the form of: i) the binary fluorosulfates of the group 4 elements, along with an investigation of their intrinsic acceptor ability as ascertained by the synthesis of Cs 2 [M(SO 3 F) 6 ] (M = Ti, Zr or HO; and ii) fluoride fluorosulfates of groups 5 and 15. Another aspect of this theme is the use of synthetic strategies such as metal oxidation by S 2 0 6 F2 in HSO 3 F, and the re-examination of the role FOSO2 F can play as a fluorinatingfluorosulfonating reagent in oxidative addition reactions. The attempts to extend an anion basicity scale for various weak nucleophiles to the carboxylates using the (CH 3 ) 2 Sn 2+ moiety as a probe also fall within this theme, as do the studies to trap Au 2+ , either in HSO 3 F solution or in solid [Au(SO 3 F)3] 2 (as lattice defects), by using highly acidic or weakly nucleophilic  166  species. The approach used in all these instances has its roots in classical synthetic inorganic fluorine chemistry. The second theme is the widespread, albeit judicious, application of modern spectroscopic techniques (FT-IR, Raman, multinuclear NMR, ESR and MOssbauer) assisted by computational methods and spectral simulations, and sometimes by single crystal X-ray diffraction, in the characterization of the highly reactive and, sometimes, corrosive systems that were developed in the pursuit of the first theme. These systems, by their very nature (physical and chemical) often require a multi-pronged approach. Here, the approach has been based on physical inorganic chemistry. It is hoped that a useful cohesion of both these themes has been achieved to a large extent, and that these efforts have given us a renewed understanding of some of the systems discussed in this thesis. As is the case with many journeys, not all the intended destinations have been reached, and neither have all the problems been solved. These constitute the bulk of the suggestions for future work.  SUMMARY AND SUGGESTIONS FOR FUTURE WORK With the synthesis of the tetrakis(fluorosulfates) of the group 4 elements, the number of known binary tetrakis(fluorosulfates) has been nearly doubled. While these compounds are not conducive to solution studies in fluorosulfuric acid, it may be worthwhile to make use of their anions to stabilize unusual cations as previously reported for Sn(SO 3 F) 4 . In recent months, the chemistry of gold has flourished considerably in our group; several carbonyl derivatives have been isolated and characterized. Based on the similar behaviour observed for platinum and palladium, it is recommended that the synthesis of compounds with these metals in unusual oxidation states be attempted in an analogous manner to that used for gold. There are indications that paramagnetic lattice defects may be generated in platinum tetrakis(fluorosulfate) and platinum tetrafluoride; previous work has also shown that, under certain circumstances, palladium may be stabilized in the +3 oxidation state.  167  The three carboxylates selected in this study indicate that the carboxylic acid derivatives are not a suitable choice for studying anion basicities. The structural diversity of the carboxylate compounds is, for this purpose, a disadvantage. From the three examples chosen in the present work, it is evident that the carboxylate group may coordinate in an unpredictable manner. However, other weak nucleophiles such as the fluoroallcylphosphonates and other hexafluorometallates are yet to be looked at, and an attempt to fit these on the correlation should be made. In the event that they do not fall on the plot, structural analyses should be carried out to ascertain whether these compounds meet the criteria imposed in this relationship. Binary fluorosulfates or fluoride fluorosulfates of high-valent molybdenum have not been isolated in the present study, although several attempts were made in this direction; tetrakis(fluorosulfato)dimolybdenum(II) remains the only binary fluorosulfate of molybdenum known so far. The addition of ligands such as acetonitrile across the quadruple Mo-Mo bond in compounds such as tetrakis(fluorosulfato)dimolybdenum(II) opens up a new avenue to novel lower valent compounds of molybdenum, and could lead to some interesting compounds. The quest for suitable pentafluoromolybdenum(VI) derivatives remains incomplete and a magnetochemical study of molybdenum pentafluoride complexes, including the hexafluoromolybdate(VI) salts, needs to be completed. The chemistry of fluorine fluorosulfate has shown that while it seems to be a very attractive reagent in principle, it is, in reality, quite uncooperative at most times. It now appears that there is no novel chemistry or unique use for this reagent to compensate for the risks taken in preparing, purifying, storing and reacting this enigmatic chemical. On a more positive note, it appears that the gas phase structures of fluorine fluorosulfate and the other halogen fluorosulfates have been solved by electron diffraction (in collaboration with Prof. H. Oberhammer in Tiibingen); it is hoped that detailed structural information may lead to a better understanding of this compound. Bearing in mind the history of this chemical and our experiences with it, it is recommended that any further chemistry involving fluorine fluorosulfate should be carried out only under the greatest caution with amounts not exceeding  168  one to two millimoles of the reagent. Larger quantities of fluorine fluorosulfate are extremely difficult to handle; they can, and will, explode with very little provocation.  CONCLUDING REMARKS Some rudimentary insights have been realized in the course of the present study, and are briefly summarized here: a)  A coordinatively unsaturated metal centre may either give rise to Lewis acidity or lead to oligomerization or polymerization. Useful conjugate superacids are formed only when the latter tendency can be fully or partly suppressed in solutions of a strong protonic acid.  b)  Fluoride and fluorosulfate ligands have many similar features and are often substituted freely with one another. The resulting oligomeric or polymeric elemental fluoridefluorosulfates may not necessarily be stoichiometric molecular compounds but rather be non-stoichiometric phases.  c)  The use of computational techniques and spectral simulations to supplement spectroscopic measurements, demonstrated in the present work for ESR and NMR, allows for better understanding of structure and bonding in compounds which are not suitable for direct structural determination by diffraction methods.  d)  Seemingly simple reactions like the oxidative addition of fluorine fluorosulfate to the sulfur fluorides or of bis(fluorosulfuryl) peroxide to molybdenum pentafluoride may not be simple, but quite often complex, and should be carried out only after carefully weighing the risks associated with using these reagents. The richness of the chemistry within the confines of inorganic fluorine chemistry is  reflected, to some extent, by the scope of this thesis. While there has often been a preoccupation with the stoichiometric nature of compounds, it is apparent, from the present study, that some of the compounds previously thought to be well-defined and molecular in nature are, in fact, non-stoichiometric and oligomeric. In instances involving mixtures of non-  169  stoichiometric compounds, microanalysis does not play a definitive role. It has been useful to wed modern spectroscopic techniques with classical preparative chemistry throughout the course of this study; while newer techniques in characterization are often encountered, it is often imperative to use well-established methods in synthetic chemistry. In many cases, one often achieves unexpected (and sometimes undesired) results, especially when working with extremely reactive reagents. It is notable that the current sentiments were summed up almost a hundred years ago by Moissan when he said, "The study of fluorine compounds is still full of surprises".  170  APPENDIX A. THE GENERATION OF GOLD(II) IONS  INTRODUCTION General Comments Gold, along with silver, platinum and palladium, belongs to the group of metals termed as the "Noble Metals". This historical classification was as a result of their inertness towards mineral acids, and their low chemical reactivities. The most prominent oxidation states of the metals in group 11 are listed in Table A.1 along with their ionization potentials.  Table A.1. Ionization potentials for group 11 elements [1] ^ 1st IP^2nd IP^3rd IP^n=1E2 IP n=1E 3 IP Ox. State observed ^ in compounds. (a)^(b)^(c)^(a+b)^(a+b+c) Cu  745  1958  3554  2703  6257  1, 2, 3, 4  Ag  731  2074  3361  2805  6166  1, 2, 3, 5  Au  890  1980  2943  2870  5813  -1, 1, 3, 5  The preferred oxidation states of the group 11 elements are underlined in the data in Table A.1; the first ionization energy is the smallest for silver, the sum of the first and second ionization energies is the smallest for copper, and the sum of the first three ionization energies is the smallest in the case of gold. This observation agrees with the fact that principal oxidation states for silver, copper and gold are 1, 2 and 3, respectively. Most of the chemistry of gold shows that the gold atom is preferentially in an oddnumbered oxidation state. Gold predominantly exhibits the +3 oxidation state and the d 8 electronic configuration gives rise to square planar coordination about the central metal in binary compounds. Examples of gold in the +5 oxidation state are only seen in AuF 5 and  171  some hexafluoroaurate(V) ([AuF6] -) complexes, where the gold atom is octahedrally coordinated; these compounds are strong oxidizing/fluorinating agents, as are most high valent fluorides. Univalent gold compounds with the heavier halogens in the form of AuX and [AuX 2] - (X = Cl, Br, I or pseudohalogen) exhibit linear coordination about gold. The nonmetallic, stoichiometric alloy CsAu is the only example of gold in the -1 oxidation state. There are fewer examples of gold compounds in which the element has an evennumbered formal oxidation state. There are no well-characterized gold(IV) complexes, and compounds with gold in the +2 oxidation state are rare [1]-[4]. The previously reported examples of compounds with gold formally in the +2 oxidation state may be classified as: a) Pseudogold(II) compounds that comprise polynuclear Au(I)/Au(III) compounds  °  which show the propensity of Au(II) (d 9 ) to disproportionate giving Au(I) (d' ) and Au(III) (d 8 ). Gold(II) chloride, AuC1 2 (Au 4 C1 8 ) [5], is a good example in which the Au(I) exhibits linear coordination and Au(III) exhibits square planar coordination (see Fig. A-1). As expected, these compounds are diamagnetic. b) Binuclear gold(II) compounds incorporating a gold-gold bond, such as Au2I 2 [/.4-(CH2)2P(CH 3 ) 2] 2 and related compounds of the type Au 2 X 2 V(CH 2) 2 P(R) 2] 2 (X = halogen, R = CH 3 , C 2 H5, t-C4H9) [4][6].  -_,....,.. \ C1 ^ Au ^ Cl Cl ^ Au  ^Cl\ -....„, Au -...„ \ CI Cl  Fig. A-1. Structure of Au[AuCI 4 ] (Ref. 5).  172  c) Mononuclear anionic or neutral complexes of gold(II) containing chelating, anionic ligands of various types. Examples of anionic complexes include [Aun(B 9 C 21.1 11 ) 2] 2[7], containing 7-bonding 1,2-dicarbollides, and [Au(1,2-S 2 C 2 (CN)2) 2] 2- , involving chelating dithiolates [8]. Neutral phthalocyanine complexes of Au(II) complete this category of compounds. It appears that the Au(II) centre is stabilized by good a-donor and 7-acceptor ligands in these mononuclear complexes. The compounds in this group are paramagnetic and have been studied extensively by ESR. In these compounds with large ligands, the unpaired electron appears to be localized predominantly in the ligandbased orbitals to give rise to radical anions. It is interesting to note that the vast majority of these gold(II) complexes have Au in a square planar environment. The principal compound and starting point in the present study is gold(III) fluorosulfate, Au(SO 3 F) 3 . Solutions of Au(SO 3 F) 3 are known to exhibit superacid behaviour [9]. Single crystals of Au(SO 3 F) 3 obtained by heating a microcrystalline sample of Au(SO 3 F) 3 in S 2 0 6 F 2 under 02 pressure have aided in the determination of the molecular structure. The crystal structure of Au(SO 3 F) 3 (Fig. 1-1, Chapter 1) shows that the compound exists as a dimer in the solid state, with two symmetrically bridging fluorosulfate groups, and two terminal fluorosulfate groups attached to each gold atom resulting in a slightly-distorted square planar environment for the gold atom [10]. Long intermolecular Au••0 contacts complete a rather distorted octahedron as seen in Fig. A-2. The original synthesis of gold(III) fluorosulfate, Au(SO 3 F) 3 [11] involves the oxidation of gold powder by bromine(I) fluorosulfate, BrSO 3 F, and proceeds via a solid intermediate, originally formulated as Au(SO 3 F) 3 .2BrSO 3 F, and subsequently identified as Br 3 [Au(SO 3 F) 4 ] [12]. This diamagnetic intermediate when pyrolyzed at elevated temperatures (ca. 100°C) gave samples of Au(SO 3 F) 3 that exhibited weak paramagnetism (2eff  =  0.54 BM, Xinc"r  follows Curie-Weiss law down to 100 K). In contrast, the alternate route to Au(SO 3 F) 3 involving the oxidation of gold powder by S 2 06 F2 in HSO 3 F [12] yields diamagnetic  173  Fig. A-2. The coordination environment of Au(III) in [Au(SO 3 F) 3 ] 2 (Ref. 10). Interatomic distances are in  A.  Au(SO3F)3 (x.corr = _147x10 -6 cm 3 •mo1 -1 ) that is isolated in vacuo from its solution in HSO 3 F at room temperature. Two possible causes for the observed paramagnetism may be ruled out: a) temperature independent paramagnetism (TIP) is unlikely since the Curie-Weiss law is obeyed between 295 and 100 K. Its presence in some samples but not in others would also imply two different structural forms for Au(SO 3 F) 3 . b) The dibromine cation, Br 2 + , so far only seen in Br2 + [Sb 3 F 16 I or in superacid solution, could conceivably be formed via the pyrolysis of the Br3 + cation in the intermediate Br 3 [Au(SO 3 F) 4 ]. The Br 2 + cation has a 2 113/2g ground state [13] and the magnetic susceptibility does obey the Curie-Weiss law. However, samples of Au(SO 3 F) 3 obtained via pyrolysis of Br 3 [Au(SO 3 F) 4] do not contain any bromine [12]. The magnetic behaviour may be explained by the presence of Au 2+ ions as defects in the Au(SO 3 F) 3 crystal lattice.  174  The pyrolysis of high valent fluorosulfates, resulting in the formation of lower valent fluorosulfates via the elimination of S0 3 F* radicals, has been reported previously for a limited  r  (SO3F)6] decomposes at 160°C to give number of compounds. The mixed-valence PdII ru IV A  Pd(SO 3 F) 2 and S 2 06 F2 [14], Ag(SO 3 F) 2 has been pyrolyzed at 215°C to quantitatively yield AgSO3F and S20 6 F2 [15], and the xenon(II) fluorosulfate, Xe(SO3F)2, decomposes at 45°C to give Xe and S 2 06 F2 [16][17]. It is therefore not unlikely that the pyrolysis of Br3 [Au(SO 3 F) 4 ] results in the reductive elimination of BrSO3 F and SO 3 F . radicals to give a material of the type Au(SO 3 F) 3 „ (0 < x < 3). This unusual magnetic behaviour has been reported for AuF 3 where a sample obtained by fluorination of gold by F2 is diamagnetic [18] while the fluorination of Au with BrF 3 followed by pyrolysis of the intermediate BrF 2 [AuF 4] at 180°C [19] produces wealdy paramagnetic AuF 3 [20]. The behaviour discussed here has also been observed for PtF 4 [20] and Pt(SO 3 F) 4 [21]. The work done in the present study was carried out in collaboration with Germaine Hwang. This appendix, in part, discusses the controlled pyrolysis of Au(SO 3 F) 3 carried out with the intent of forming Au 2+ cations as lattice defects in Au(SO 3 F) 3 , and the subsequent identification of these cations by ESR. This route has some advantages over the pyrolysis of Br3 [Au(SO 3 F) 4] because: a) Vibrational spectra on Au(SO 3 F) 3 have been reported [12] and may be obtained without much difficulty. b) It is hoped that knowledge of the structure of Au(SO 3 F) 3 will simplify the interpretation of the ESR measurements. c) The anticipated elimination of SO 3 F . or S 2 06 F2 is straightforward, and the reaction may be monitored by a pressure increase and a weight loss in the material. d) The pyrolysis of Br 3 [Au(SO 3 F) 4 ] has yielded samples of Au(SO3F)3 that give rise to broad ESR signals, possibly due to a high concentration of paramagnetic centres in the sample. Estimates on samples used for magnetics suggest ca. 7% of all the gold is present as Au 2+ . Another potential route to Au 2+ ions arises from some work done recently in our group. The weakly basic tetrakis(fluorosulfato)aurate(III) anion, [Au(SO 3 F) 4 ]', present in the monoprotonic superacid system, HSO 3 F-Au(SO 3 F) 3 , is capable of stabilizing strongly  175  electrophilic cations [12]. The strong acidity of this system and the apparent lack of oxidizing power of Au(III) has motivated attempts to protonate carbon monoxide in the hope of obtaining the formyl cation, HCO + . This experiment has resulted in the formation of lowvalent carbonylgold(I) species such as monocarbonylgold(I) fluorosulfate, Au(CO)SO 3 F [22] and [Au(CO)2] +, with CO acting as a reducing agent and coordinating ligand. It is this reducibility of Au(III) in HSO 3 F that has prompted the present attempt to obtain Au 2+ ions in solution of HSO 3 F by the reduction of Au(SO 3 F) 3 , under mild conditions using gold powder. The solution behaviour of Au(SO 3 F) 3 in HSO3F is well known [9][12][23]. This section, in part, discusses the generation of solvated Au 2+ in solutions of HSO 3 F by either dissolving pyrolyzed Au(SO 3 F) 3 or reducing gold(III) by a suitable reagent. The Au 2+ ions in HSO 3 F solution are identified by ESR spectroscopy. It should be noted that the stabilities in the oxidation states of the group 11 elements are of practical importance. For example, although the mechanism of high-temperature superconductivity in the well-known superconductor La 2 _xSrxCu04 [24] is not clear, it is believed that copper atoms in the +2 oxidation state may play a key role [25]. The lack of any high-temperature superconductors of silver and gold is attributed to the instability of the +2 oxidation state of these elements.  ESR Spectroscopy in the study of Au 2+ ions Electron spin resonance is one of the few techniques that can be used to study free radicals. It can also be used to study small amounts of paramagnetic species trapped in a diamagnetic matrix. The observation of an ESR spectrum from a sample is an indication that at last a small portion of the sample has unpaired electrons. It is useful for the study of unstable transient paramagnetic species that may be generated in situ by methods such as electrochemical oxidation-reduction or pyrolysis as is in the present study. Since low concentrations of Au 2+ are expected, it is not necessary to dilute the samples with an isomorphous diamagnetic compound.  176  ESR spectroscopy is based on the spin of an electron and its associated magnetic moment. In the presence of a magnetic field, B o , a molecule (or ion) having one unpaired electron has two electron spin energy levels given by:  E = gp, B Bo Ms where AB = Bohr magneton  Ms = electron spin quantum number +1/2 (for a doublet) g = proportionality factor = 2.00232 for a free electron  For organic free radicals, the g value is very close to that for a free electron, but when dealing with transition metals, the spin-orbit coupling and zero-field splitting can lead to pronounced variations in g values. Transitions between such energy levels are studied by ESR. These transitions are induced by an application of radiation of appropriate frequency. When B o = 0.34 T, frequencies of ca. 9.5 GHz in the microwave region are used. When the electron interacts with a neighbouring nuclear magnetic dipole, these energies then become: E = gA B B oM s + AM s m i where A = hyperfine splitting m 1 = nuclear spin quantum number for the neighbouring nucleus  The appearance of hyperfine coupling provides considerable information about the species involved. The coupling may arise in two ways: (i) direct dipole-dipole interaction, which depends on the angle between the vector joining the two dipoles and the applied magnetic field, and (ii) Fermi contact interaction, which is due to unpaired electron density at  177  the nucleus in question. Fermi contact interaction is only present when the electron resides in an orbital with some s character. Observation of coupling in the latter case gives direct indication of the extent of electron delocalization. In anisotropic systems, which we encounter when unstable species are trapped in polycrystalline media, solids, frozen solutions, and paramagnetic point defects in single crystals, the g factors may be considered an asymmetric tensor. When diagonalized, this tensor gives us the three principal g factors, gam , gyy and gzz . In systems with spherical or cubic symmetry, they are equal ie. g  ^gyy = gzz = giso ; in solutions, they are averaged to  giro. In systems with axial symmetry (such as square planar sites) we have g zz = g o g  ^gyy = g_i_ and giro = (g o + 2 gam) / 3  In systems with lower symmetry we have three distinct g factors, g ^gyy^gzz ; giso is the average. Similarly, the hyperfine splitting parameter, A, must be replaced by A, a , Ayy and A u , the same arguments as above being applied to get A xx = A yy = A_L and A u = go for symmetric tops and all these terms being different for systems with lower symmetry. Hence ESR spectroscopy serves as a useful probe into Au(II) systems and may provide some indication of where the unpaired electron resides in the complex.  EXPERIMENTAL Reagents Gold powder (99.95 purity, -20 mesh) was obtained from the Ventron Corporation (Alfa Inorganics). The fluorosulfuric acid and bis(fluorosulfuryl) peroxide were obtained and purified as outlined in Chapter 2. Gold tris(fluorosulfate), Au(SO 3 F) 3 , was prepared according to published methods [121  178  Synthetic Reactions These are covered in detail elsewhere [26] and are reported here to show how the Au 2+ anions may be generated.  Pyrolysis of Au(S0 3 E13. 265 mg (0.536 mmol) of Au(SO 3 F) 3 were heated in a 100 mL round bottom reaction vessel at 60°C for 7 days in a static vacuum. There was no discernable color change or sublimate. Although there were no volatile products at -196°C, about 9.8 Amol of volatiles were obtained and removed in vacuo at room temperature. The compound was then heated to 100°C for ca. 4 days, and an orange sublimate was visible on the walls of the reactor. Upon cooling to room temperature, about 7 Amol of volatiles were measured and removed in vacuo. The compound was then heated to 125°C for ca. 2 days upon which about 5.8 Amol of volatiles were measured and removed in vacuo. The sublimate was then scraped and the heating was continued for another 7 days. At this point, the sample had decreased in weight by 17.6 mg (0.036 mmol), and the sulfur content had dropped from 19.5% to 18.9%. The sample was transferred to 1 mm capillary tubes and ESR spectra were obtained as described in Chapter 2. Reduction of Au(III) to Au(II) in HSO 3 F by Au. A mixture of 255 mg (0.516 mmol) of Au(SO 3 F) 3 and 42 mg (0.213 mmol) of Au metal was dissolved in about 3 mL of HSO 3 F and stirred at room temperature. The initial solution was opaque orange in color, but after about 15 minutes, the color changed to orange-red. The turbid solution eventually became clear after about 4 hours of stirring at room temperature. The reaction mixture was heated overnight at 65°C upon which all the Au metal had dissolved, and the solution had a dark orange-brown color. After three days of stirring at 25°C, a yellow-orange solid was formed. The solid was isolated by vacuum filtration and analyzed as Au(SO 3 F) 2 ; the solid was "ESR silent". The color of the solution remained unchanged and a low temperature ESR spectrum of the solution was recorded.  179  DISCUSSION The pyrolysis of Au(SO3F)3 was followed by weight determination. The total weight loss was about 6% and the decrease in sulfur content was about 3%. This observation suggests that the volatiles, discarded after each temperature increase, are sulfur containing species. The infrared spectrum of the product showed subtle changes from that of Au(SO 3 F) 3 before pyrolysis, primarily in the SO 3 F-stretching region. These changes in the infrared spectrum are attributed to the loss of S 2 06 F 2 leading to the formation of Au +2 cations present amongst the Au(III) sites. The infrared frequencies of partly pyrolyzed Au(SO3F)3 are comparable with those of Au(SO 3 F) 3 and Au(I)Au(III)(SO 3 F) 4 . The differences in the spectra are rather subtle, mainly in the SO 3 F stretching region and the Au04 skeletal vibrations. The spectrum of Au(SO 3 F) 3 shows greater band proliferation after pyrolysis, and the band at the highest frequency (1442 cm -1 ) has shifted to lower frequency (1434 cm -1 ) and has decreased in intensity. Although the infrared spectrum of Au(I)Au(III)(SO3F)4 is slightly different, it is possible that this compound may be present in small quantities in the pyrolyzed Au(SO 3 F) 3 . Hence, the infrared data indicate that, upon heating, Au(SO 3 F) 3 undergoes subtle structural changes that increase the complexity in the infrared spectra. It is proposed that the distorted square planar coordination present in Au(SO3F)3 is maintained by a proportionally greater number of (originally) terminal SO 3 F groups now functioning as bidentate bridging SO 3 F groups. The observed weak paramagnetism of pyrolyzed Au(SO 3 F) 3 may be explained by partial reductive decomposition according to: Au(SO3F)3 ^ > A U ( S 03F )3_x + x/2 S 2 06 F2 The pyrolysis of AuF 3 is reported to produce F2 and gold in an analogous manner [19]. The ESR spectrum of Au(SO 3 F) 3 x , obtained at 100 K (Fig. A-3a), shows the presence of a radical with approximately axial symmetry (g 1 = 2.093, g 2 = 2.103, g 3 = 2.882, and Lorentzian linewidth = 32 G) and an unusually high  giro  value of 2.359; there is no resolved  hyperfine coupling discernable. A small isotropic feature (g = 2.0033) was occasionally  Lower spectrum simulated with:  250 G  = 2.103 giy. ---- 2.093 gzz = 2.882 Lorentzian Linewidth = 32 G Fig. A-3. X-band ESR spectrum of Au(SO 3 F) 3 _x recorded at 100 K  B  181  observed in addition and, based on g value comparison [27]-[29], is attributed to the SO 3 F  .  radical, presumably trapped in the sample. In order to study the volatile decomposition products during the pyrolysis of Au(SO 3 F) 3 , two shorter (ca. 24 h) experiments were carried out at 115 and 145°C. The major products identified by infrared spectroscopy were sulfuryl fluoride (SO 2F2) and silicon tetrafluoride (SiF4) with S 2 06 F2 and S205F 2 as the minor constituents. The formation of SO 2F 2 may be explained by a decomposition mode of fluorosulfates that has been observed for Ba and Sr [30][31]:  500°C M(SO 3 F) 2 ^ > MSO4 + SO2F2 ^ -2M = Ba or Sr  However, the temperatures required for this decomposition to take place are considerably higher than those encountered in the present study, and no sulfate species is detected by infrared spectroscopy. An alternative explanation that also accounts for the observation of SiF 4 involves the reaction of SO 3 F radicals with the Pyrex walls of the reactor. The reaction .  of SO 3 F radicals with quartz above 120°C has been reported to give SO 2 F2 as well as .  SiF 4 [32]. This interpretation is supported by the presence of trapped SO 3 F radicals in the .  solid pyrolyzed residue (vide supra). In addition, the sublimate in these experiments was found to be "ESR silent", whereas the sublimation of Au(SO 3 F) 3 in a dynamic vacuum also yielded a product that gave rise to an ESR signal. Identical ESR spectra were obtained from samples of Au(SO 3 F) 3 that were sublimed in  vacua. ESR spectra obtained on Cs[Au(SO 3 F) 4] heated briefly to 100°C and Cs[Au 2 (SO 3 F) 7] melted at 102°C [33] are also identical to those obtained in this work. In contrast, gold(III) fluorosulfate derivatives which are not heated do not give rise to any signals detectable by ESR. Furthermore, when the partly pyrolyzed "ESR active" samples are treated with S206F2  182  they become "ESR silent", and a small gain in weight is noted. In all these experiments, gold(II) appears to be in a nearly square planar site. A differential thermal gravimetry experiment, courtesy of Prof. Dr. Meyer (Hannover, Germany),indicates that the Au(SO3F)3 decomposes gradually between 130-350°C. Distinct decomposition steps are not detected and some sublimation does occur. The overall weight loss is a little greater than 50%. This indicates that a stepwise pyrolysis of Au(SO 3 F) 3 to give Au(SO 3 F) 2 or AuSO 3 F does not occur at any point. When gold powder and Au(SO 3 F) 3 were reacted at 65°C in HSO3F, all the gold dissolved after several days and a yellow-brown solid precipitated from a dark brown solution. The isolated precipitate was analyzed as the previously unknown gold(II) fluorosulfate, formed according to:  HSO3F 2 Au(SO3F) 3 + Au ^ > 3 Au(SO3F)2^ -3-  The diamagnetic compound should be viewed as a pseudogold(II) complex and is better formulated as Au(I)Au(III)(SO 3 F) 4 . The ESR spectrum of the supernatent brown solution is totally asymmetric when the ratio of Au(SO 3 F) 3 :Au is exactly 2:1. When this ratio is  ca. 2.5:1, with some excess Au(SO 3 F) 3 remaining in solution after completion of the reduction, the ESR spectrum exhibits axial symmetry. A frozen sample of a solution with Au(SO 3 F) 3 in excess gives rise to the ESR spectrum shown in Fig. A-4a. A disproportionation reaction of the type: 2 Au 2+ (solv) + 4 SO 3 F -  Au(I)Au(III)(SO3F)4^-4-  is suggested. The near-axial spectrum obtained at 100 K from the centre produced by the reduction of Au(SO3F)3 with Au in HSO3F exhibits hyperfine splitting, most probably due to  197 Au,  an  I = 3/2 nucleus. It was necessary to include a nuclear quadrupole coupling in the simulation  P  Lower spectrum simulated with: A 145 MHz An, = 145 MHz A. = 120 MHz Q.= 46 MHz  250 G  gu = 2.103 gn, = 2.096 g.= 2.890  Gaussian Linewidth = 32 G Fig. A-4. X-band ESR spectrum of Air(sole) recorded at 100 K  B  I  184  in order to obtain the correct relative line intensities in the spectrum as seen in the g - 2.1 region; this confirms the fact that the paramagnetic centre is indeed Au-based. The parameters used in the simulation of this spectrum (g 1 = 2.093, g2 = 2.103, g 3 = 2.890, A l = A2 = 145 MHz, A3 = 120 MHz and Qzz = 46 MHz; Gaussian linewidths L 1 = L2 = 20 G  and L3 = 25 G) indicate that the g tensor is the same as that of the species obtained from the pyrolysis of Au(SO3F)3, within experimental error. The orthorhombic nature of the g tensor for both Au(II) centres suggests that, in both instances, the Au(II) occupies a site very similar to the the Au(II) site formed as a lattice defect in Au(III) fluorosulfate. The observed hyperfine splitting of four lines in the perpendicular region of the ESR spectrum of the solvated samples is also consistent with the electron localized on 197 Au (I = 3/2). The absence of hyperfine coupling in Au(SO 3 F) 3 _x is attributed to exchange narrowing in the solid state due to electron interactions in the oligomers formed by the loss of S 2 06 F 2 . This is evident from the fact that although  giro  values are identical for both species, a  greater intrinsic line width is needed to simulate the spectrum of Au(SO3F)3 x , the gold centre formed via partial pyrolysis of Au(SO 3 F) 3 . The nature of this motional averaging, however, is open to speculation. In both the instances that Au(II) is detected, there is some Au(III) present as well. Au(SO 3 F) 3 is dimeric in the solid state [10], and is believed to be dimeric in HSO 3 F as well [23]. It is, therefore, likely that the Au(II) centre is found in an oligomeric structure wherein Au(III) strongly influences the overall structure and is responsible for the square planar environment of both the Au(III) and Au(II) centres. When diamagnetic Au(I)Au(III)(SO 3 F) 4 is re-dissolved in HSO 3 F, an asymmetrically coordinated Au(II) species is formed in solution and is detected by ESR spectroscopy. This is similar to the Au 2+ sol v generated from a 2:1 mole ratio of Au(SO 3 F) 3 :Au in HSO 3 F. The giro value of this spectrum is 2.318; the presence of a monomeric Au 2+ species is postulated. Addition of Au(SO 3 F) 3 to this solution once again gives a product with near-axial symmetry.  185  The large g shifts (g iro = 2.362) from the free spin value suggest that the unpaired electron is largely localized on a species with high spin-orbit coupling, consistent with Au(II) (ca. 5000 cm -1 ). Silver(II) fluorosulfate derivatives exhibit similar g anisotropy and high giro  values, for example, Ag(SO 3 F) 2 has giro = 2.220 reflecting a smaller spin-orbit coupling constant for Ag 2+ .  CONCLUSIONS The solvent used in this study, HSO 3 F, is one of the strongest protonic acids and has been known to stabilize unusual cations; its self-ionization anion SO 3 F - is one of the least basic anions. Neither is susceptible to reduction and neither is capable of acting as a 7-acceptor unlike the ligands used previously [4][6]. The ESR spectra obtained in this work, therefore, differ dramatically from those previously reported for Au(II) complexes [2]-[8][34]. The g anisotropy is very much greater and the giro values are substantially higher than previously published values. This strongly suggests the formation of centres in which the odd electron resides principally on the metal rather than on the ligands as is argued previously [2]-[8]. The low stability of Au 2+ towards disproportionation is borne out in the fact that pure paramagnetic Au(SO 3 F) 2 was not obtainable in the present study, and that few examples of Au 2+ are known to exist. All the experimental evidence presented here, however, convincingly points to the formation of Au 2+ ions stabilized by very wealcly basic fluorosulfate groups.  186  REFERENCES  1. R.J. Puddephatt, in G. Wilkinson (ed.), "Comprehensive Coordination Chemistry", Vol.7, Pergammon, Oxford (1987) p. 861. 2. R.J. Puddephatt, "The Chemistry of Gold", Elsevier, Amsterdam (1978). 3. B.J. Muller, Angew. Chem. Int. Ed. Engl., 26 (1987) 1081. 4. H. Schmidbauer and K.C. Dash, Adv. Inorg. Chem. Radiochem., 25 (1982) 239. 5. D.B. Dell'Amico, F. Calderazzo, F. Marchetti, S. Merlino and G. Perego, J. Chem. Soc.  Chem. Comm., (1977) 31. 6. J.P. Fackler and J.D. Basil, in M.H. Chisolm (ed.), "Inorganic Chemistry Towards the 21st Century", ACS Symposium Series, 211, American Chemical Society, Washington (1983) p. 201; Organometallics, 1 (1982) 871. 7. L.F. Warren and M.F. Hawthorne, J. Am. Chem. Soc., 90 (1968) 4823. 8. R.L. Schlupp and A.H. Maki, Inorg. Chem., 13 (1974) 44. 9. K.C. Lee and F. Aubke, Inorg. Chem., 18 (1979) 389. 10. H. Willner, S.J. Rettig, J. Trotter and F. Aubke, Can. J. Chem., 69 (1991) 391. 11. W.M. Johnson, R. Dev and G.H. Cady, Inorg. Chem., 11 (1972) 2260. 12. K.C. Lee and F. Aubke, Inorg. Chem. 19 (1980) 119. 13. a) W.W. Wilson, R.C. Thompson and F. Aubke, Inorg. Chem., 19 (1980) 1489. b) M.S.R. Cader, R.C. Thompson and F. Aubke, Can. J. Chem., 67 (1989) 1942. 14. K.C. Lee and F. Aubke, Can. J. Chem., 55 (1977) 2473. 15. P.C. Leung and F. Aubke, Inorg. 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Soc., 91 (1969) 5936. 29. W.V. Cicha, F.G. Herring and F. Aubke, Can. J. Chem., 68 (1990) 102. 30. W. Traube, J. Hoerenz and F. Wunderlich, Ber., 52 (1919) 1272. 31. E.L. Muetterties and D.D. Coffman, J. Am. Chem. Soc., 80 (1958) 5914. 32. F.B. Dudley and G.H. Cady, J. Am. Chem. Soc., 85 (1963) 2375. 33. G. Hwang, B.Sc. Thesis, University of British Columbia, Vancouver, 1990. 34. Landolt-BOrnstein. Numerical data and functional relationships in science and technology. Vol 2. Magnetic properties of coordination and organometallic transition metal compounds. Springer-Verlag, Berlin (1966).  


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