Fluorosulfates of Silver, Ruthenium, and Osmium

FLUOROSULFATES OF SILVER, RUTHENIUM, AND OSMIUM

PATRICK CHEUNG SHING LEUNG

B.Sc. (Hons.), The University of British Columbia, 1975

A THESIS SUBMITTED IN PARTIAL FULFILMENT OF

THE REQUIREMENTS FOR THE DEGREE OF

DOCTOR OF PHILOSOPHY

THE FACULTY OF GRADUATE STUDIES

(Department of Chemistry)

We accept this thesis as conforming

to the required standard

THE UNIVERSITY OF BRITISH COLUMBIA

December, 1979

© Patrick Cheung Shing Leung, 1979 In presenting this thesis in partial fulfilment of the requirements f an advanced degree at the University of British Columbia, I agree tha the Library shall make it freely available for reference and study.

I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the Head of my Department or by his representatives. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission.

Department of CMMUTtY

The University of British Columbia 2075 Wesbrook Place Vancouver, Canada V6T 1W5 ABSTRACT

A number of synthetic routes to silver(II) fluorosulfate,

Ag(SC>3F)2, were systematically explored. The most suitable and versatile route was found to be the oxidation of silver

metal by a solution of bisfluorosulfuryl peroxide, S20gF2, in fluorosulfuric acid, HSO^F, according to:

HS03F

Ag + S206F2 m- Ag (SC>3F) 2>

Additional methods which were found to be suitable involved the oxidation of a wide variety of silver(I) compounds such as

S F or the nsert:LOn of S0 Ag20 or AgSO^F by 2°6 2' i- 3 into AgF2.

Structural conclusions on Ag(S03F)2 and the other compounds synthesized subsequently were based on the vibrational, elec• tronic and electron spin resonance spectra, as well as on magnetic susceptibility measurements made between 300 and 77 K.

Ag(S0.jF)2 was found to be a true compound of divalent silver, 2 + with the Ag ions in either square planar or tetragonally elongated octahedral environment. The only other example of

a binary silver(II) compound is AgF2.

Several silver(II) fluorosulfate derivatives were prepared

and characterized. The reactions of bromine(I) fluorosulfate

with metallic silver and other silver(I) substrates resulted

in a mixed valence complex Ag^Ag''"'''(SO^F)4 . Its potassium analogue r^AgCSO^F)^, as well as two hexakisfluorosulfato- metallate (IV) complexes AgPtIV (SO..F) , and AgSnIV (SO-.F) . , and 3 b j b the N-donor ligand complex [Agtbipy^l (SO^F^ (i-n which bipy = 2,2'-bipyridine) were also synthesized. The attempt

to synthesize a silver(III) fluorosulfato complex by direct

insertion of SO^ into CsAgF^ resulted in the fluorination

S F and of SC>3, to give 2°6 2 CsAg(S03F)3. Finally, the solvolysi

of Ag(S03F)2 in trifluoromethylsulfuric acid, HSO^CF^, allowed

its conversion into AgtSO^CF^),,.

The principal synthetic route, the oxidation of metal

by S2°6F2 s°lur-i°ns in HSO^F, was found to be useful in the

preparation of RutSO^F)^. Ruthenium was also found to form

a number of anionic derivatives with the metal in the +3 or

+ + +4 oxidation state, as in M[Ru(SC>3F)4] with M = Cs , C102 ;

+ + M2[Ru(S03F)6J with M = Cs , K ; and Cs[Ru(S03F)g].

Two different forms of Os(S03F)3 were found. Initial

oxidation of osmium metal with S2°gF2 yielded the bright green

a-Os(S03F)3, which was converted to the light green 3-form

on long standing in S9OfiF;?. - iv -

TABLE OF CONTENTS

Page

ABSTRACT ii

LIST OF TABLES xii

LIST OF FIGURES xvi

ACKNOWLEDGEMENTS xviii

I. GENERAL INTRODUCTION 1

A. Introductory Remarks 1

B. The Fluorosulfate Radical And Its Anion.... 6

C. Synthetic Routes To Fluorosulfates 19

D. Vibrational Studies Of The Fluorosulfate

Group 2 5

E. Transition Metal Fluorosulf ates 30

F. Structural Characterizations Of Transition

Metal Fluorosulf ates 34

II. - EXPERIMENTAL 35

A. Apparatus 35

1. Pyrex vacuum line 35

2. Metal vacuum line. 36

3. Reactions vessels 37

4. HSO-jF distillation apparatus 44

5. HSO^CF^ distillation apparatus 44

6. Drybox 4 6 - v -

Page

7. Miscellaneous 46 B. Instrumental Methods 47

1. -Infrared spectroscopy 47

2. Raman spectroscopy 50

3. Visible and ultraviolet spectroscopy.... 50

4. Magnetochemistry 51

5. Electron spin resonance spectroscopy.... 53

6. Mossbauer spectroscopy 53

7. Melting points 54

C. Chemicals 54

1. Commercial sources 54

2. Preparative reactions 57

D. Chemical Analyses 61

1. Elemental analyses 61

2. Oxidation states determinations 6 3

III. FLUOROSULFATES OF SILVER(II) 64

A. Introduction 64

B. Silver (II) Fluorosulfate 76

1. Introduction 76

2. Synthetic reactions 77

(a) Reactions with fluorosulfuryl peroxide 77

(b) Reactions with mixtures of HSO^ and

S2°6F284 - vi -

Page

and sulfur trioxide 85

(d) Reactions with bromine monofluoro-

sulfate 86

(e) Discussion 88

3. Characterization of Ag(S03F)2 96

(a) Infrared spectra 96

(b) Electronic spectra 99

(d) Electron spin resonance spectra Ill

(e) Oxidation state determination 117

(f) Thermal decomposition 118

C. Coordination Complex Of Ag(S03F)2,

[Ag(bipy)2] (S03F)2 119

1. Introduction 119

2. Synthesis of silver(II) bis(2,2'-bipy-

ridyl) bis (fluorosulfate) 120

3. Characterizations 121

(a) Vibrational spectra 121

(b) Electronic spectra 126

(d) E.S.R. spectra 130

D. Anionic Fluorosulfato Complexes Of Silver(II) 132

1. Introduction 132 - vii -

Page

2. Syntheses and elemental analyses 132

(a) Ag3(S03F)4 . 132

(b) K2[Ag(S03F)4] 132

3. Experimental results and discussions.... 135

(a) Infrared spectra 135

(b) Magnetic susceptibility measurements 135

(d) Oxidation state determination and

thermal decomposition of Ag3(S03F)4. 147

E. Silver(II) Hexakis(fluorosulfato)metallate(IV) 148

1. Introduction 148

2. Syntheses and elemental analyses 149

(a) AgPt(S03F)6 149

(b) AgSn(S03F)6 150

3. Characterizations 151

(a) Vibrational spectra . 151

(b) Electronic spectra 154

(d) E.S.R. spectra 160 119

(e) Sn Mossbauer spectra of AgSn(S03F)g 162

F. Reactions Of Ag(S03F)2 164

1. Reaction of Ag(S03F)2 with elemental

fluorine 164 - viii -

Page

(a) Introduction 164

(b) Conversion of Ag(S03F)2 into AgF2«.. 165

2. Reactions with some other halogens 167

(a) Chlorine 167

(b) Bromine 167

3. Reactions with chloryl fluorosulfate.... 169

4. Reaction with pyridine 171

5. Reaction with acetonitrile 172

6. Reaction with antimony pentafluoride.... 173

G. Attempts to obtain higher oxidation states

of silver 175

S F 175 1. The reaction of AgO and 2°6 2

2. The reaction of CsAgF4 and S03 176

(a) Introduction 176

(b) Synthetic reactions 177

IV. SILVER (II) TRIFLUOROMETHYLSULFATE 18 6

A. Introduction 18 6

B. Synthetic Reactions 188

1. Synthesis of silver(II) trifluoromethyl-

sulfate 188

(a) Reaction of Ag(SO-jF)0 and HSO^CF-,... 188 - ix -

Page

(b) Other synthetic attempts 189

2. Conversion of Ag(S03CF3)2 into

[Ag(bipy)2] (S03CF3)2 193

C. Experimental Result And Discussion 193

1. Infrared spectrum 193

2. Magnetic measurements 194

3. E.S.R. spectrum 198

4. Thermal decomposition 200

D. Conclusion 202

V. FLUOROSULFATES OF RUTHENIUM 203

A. Introduction 203

B. Ruthenium (III) Fluorosulfate 207

1. Preparation and elemental analysis 2 07

2. Experimental result and discussion .... 208

(a) Infrared spectra 208

(b) Magnetic susceptibilities 210

(d) E.S.R. spectra 217

(e) Discussion 219

C. Anionic Fluorosulfato Complexes Of

Ruthenium (III) 222 - x -

Page

1. Introduction 222

2. Syntheses and elemental analyses 222

(a) Cl02[Ru(S03F)4] 222

(b) Cs[Ru(S03F)4] 223

Cs3 [Ru(S03F) ] 223

3. Characterizations 224

(a) Infrared spectra 224

(b) Magnetic susceptibilities 227

D. Anionic Fluorosulfato Complexes of

Ruthenium(IV) 229

1. Introduction 229

2. Preparations and elemental analyses 2 31

(a) K2 [Ru(S03F) ] 231

(b) Cs2 [Ru(S03F) ] 231

3. Characterizations 232

(a) Vibrational spectra 232

(b) Magnetic susceptibility measurements 233

E. Other Synthetic Attempts 239

l.i. Reaction of S20gF2 with ruthenium metal. 239

2. Reaction of So0_F_ with Ru(SO_,F)_ 240 2 6 2 3 3

3. Reaction of BrS0oF with ruthenium metal. 24 0 - xi -

Page

4. Reactions of Ru02 241

VI. OSMIUM (III) FLUOROSULFATE 242

A. Introduction 242

B. Synthetic Reactions And Elemental Analyses

1. Synthesis of Os (S03F) 245

2. Other synthetic attempts 246

3. Attempts to synthesize fluorosulfato

complexes of osmium 247

C. Experimental Result And Discussion 248

1. Vibrational spectra 248

2. Magnetic susceptibility measurements.... 250

3. Electronic spectra 253

4. E.S.R. spectra 255

5. Conclusion 256

VII. GENERAL CONCLUSION 257

A. Summary 25 7

B. Suggestions For Further Work 258

BIBLIOGRAPHY 260

APPENDICES 277 - xii -

LIST OF TABLES

Table Page

1. Fundamental Frequencies Of Matrix Isoated

SO-jF and Comparison With Other Results 9

2. Ligand Field Parameters For Some Nickel Salts... 16

3. Selected Parameters For F~ And S03F~ 18

S F and 4. Some Physical Properties Of HS03F, 2°6 2

BrS03F 24

5. Structural Bonding Differentiation For The S03F

Group 2 6

6. Commercially Available Types Of Apparatus 48

7. Chemicals 55

8. Some Ionization Potentials Of Silver And Copper. 65

9. Synthetic Routes To Ag(S03F)2 92

10. Infrared Spectra Of Ag(S03F)2 And Related

Compounds 98

11. Electronic Spectra Of Ag(S03F)2 And Related

Compounds 102

12. Magnetic Susceptibilities And Magnetic Moments

Of Ag (S03F) 106

13. Magnetic Properties Of Ag(S03F)2 And Related

Compounds * 110 - xiii -

Table Page

14. ESR Data Of Silver (II) Fluorosulfate 116

15. Vibrational Spectra Of [Ag(bipy)2](S03F)

And [Ag (bipy)2 ] (CF3S03F) 2 123

16. Anionic Infrared Bands Of [Ag(bipy)2](S03F)

And Related Compounds 12 5

17. Electronic Spectra Of [Ag(bipy)2](S03F) And

Related Compounds 127

18. Magnetic Susceptibilities And Magnetic Moments

Of [Ag(bipy)2 ] (S03F)2 129

19. ESR Data Of Some 2,2'-bipyridine Complexes Of

Silver (II) 131

20. Infrared Spectra Of Ag3(S03F)4 and K2Ag(S03F)4.. 136

21. Magnetic Susceptibilities And Magnetic Moments

Of Ag3(S03F)4 And K2Ag(S03F)4 138

22. Experimental J Values Of Ag3(S03F)4 143

23. Experimental J Values Of K2Ag(S03F)4 144

24. Vibrational Spectra Of AgPt (S0oF) and 3 D

AgSn(S03F)6 And Related Compounds.... 153

25. Electronic Spectra Of AgPt(S03F)g, AgSn(S03F)g

And Related Compounds 157

26. Magnetic Susceptibilities And Magnetic Moments

Of AgPt(SO-F)c And AgSn (SO..F) 159

27. ESR Data Of AgPt(S03F)g And AgSn(S03F)6 And

Related Compounds 161 - xiv -

Table Page

28. Infrared Spectra of CsAg(S03F)3 and Related

Compounds » . . . . 181

29. Magnetic Susceptibilities and Magnetic Moments

for CsAg (S03F) 3 183

30. Infrared Frequencies of Ag(S03CF3)2 and Related

Compounds.. 195

31. Magnetic Susceptibility Data of Ag(S03CF3)2 197

32. Experimental J values of Ag(SC>3CF3)2 199

33. Infrared Frequencies of Volatiles from Thermal

Decomposition of Ag(S03CF3)2 201

34. Infrared Frequencies of Ru(SC>3F)3 and Related

Compounds 20 9

35. Magnetic Susceptibility Data of Ru(SC>3F)3 212

36. Electronic Spectra of Ru(S03F)3 and Related

Complexes 216

37. ESR Data of Ru(S03F)3 218

38. IR Frequencies of Anionic Fluorosulfato Complexes

of Ruthenium (III) 225

39. Magnetic Susceptibility Data of Cs[Ru(S03F)4].... 228

40. Vibrational Frequencies of Anionic Fluorosulfato

Complexes of Ruthenium(IV) 2 34

41. Magnetic Susceptibility Data of K2[Ru(S03F) ],

Cs2[Ru(S03F)g], Cs[Ru(S03F)5] 236 - XV -

Table . Page

42. Electronic Spectra of Fluorosulfato Complexes

of Ruthenium 238

43. Vibrational Spectra of Os(S03F)3 249

44. Magnetic Susceptibility Data of Os(S03F)3 252

45. Electronic Spectra of Os(S03F)3 and Related

Compounds 254 - xvi -

LIST OF FIGURES

Figure Page

1. Correlation Diagram for the Fluorosulfate group 27

2. Existing Binary Transition Metal Fluorides and

Fluorosulfates 31

3. Pyrex Reaction Vessels 38

4. Monel Metal 2-Part Reaction Vessel 42

5. Kel-F Reaction Trap 4 3

6. Fluorosulfuric Acid Distillation Apparatus 45

7. Apparatus for the Preparation of S20gF258

8. Metal Flow Reactor for Fluorination Reactions.. 62

9. Electronic Configuration and Spectroscopic

Terms of Silver(II) of Octahedral and Elongated

Tetragonal Ligand Field 101

10. Magnetic Susceptibilities of Ag(S03F)2 and

[Ag(bipy)2] (S03F)2 108

11. ESR-Spectra of Ag(S03F)2 at 80 K 115

12. The Infrared Spectrum of [Ag(bipy)2](S03F)2.... 122

13. Magnetic Susceptibility of Ag3(S03F)4 from 80

to 310 K 139

14. V^ff versus kT/J plot for A93(S03F)4142

u versus 15. eff -kT/J Plot for K2Ag(S03F)4 146

16. The Raman Spectrum of AgPt(S0-.F)fi 152 - xvii -

Figure Page

17. Diffuse Reflectance Spectrum of AgSn(SO^F)^.... 155

18. Magnetic Susceptibilities of AgPtCSO^F) and

AgSn(S03F)6 158

19. Interconnecting Reaction Scheme between the

Fluorosulf ates of Silver 170

20. The I.R. Spectrum of Cs[Ag(S03F)3]in the BaF2

Region 180

21. UV-VISIBLE Spectrum of Ru(S03F)3 in HSC>3F 214

22. The Structure of Ru2 (OCOC3H7)4 C1 220

23. Molecular Orbital Diagram for Ru2(02CC3H7)4C1.. 220

24. The I.R. Spectrum of Cs [Ru (S03F)4 ] 226

25. The E.S.R. Spectrum of Cs[Ru(S03F)4] at 80 K... 230

26. The Raman Spectrum of K2[Ru(S03F)g] 235

27. I.R. Spectra of a- and B-Os(SO,F) 250 - xvin -

ACKNOWLEDGEMENTS

I am extremely grateful to my research director, Professo

F. Aubke, for his constant advice and encouragement, both academic and personal, throughout the entire course of this work.

Special thanks are due to Professor R.C. Thompson for his many helpful comments and discussions. Thanks are also extended to various coworkers in Lab. 457 for making my years

in graduate study more enjoyable, and in particular, to

Mr. Keith Lee for his stimulating discussions and assistance

in various aspects of this work. Professors N.L. Paddock and

R.C. Thompson are thanked for reading portions of the manuscri

Finally, I would like to thank my wife, Marianne, for her typing of this manuscript, and above all, for her kind

understanding and patience. TO

MY PARENTS 1

I. GENERAL INTRODUCTION

A. INTRODUCTORY REMARKS

The transition elements may be defined as having partly filled d or f orbitals in their common oxidation states.

The large number of transition metals can be divided into three main groups: (1) d-block elements, (2) lanthanides and (3) actinides. Only the d-block elements, where often a further distinction into 3d, 4d or 5d elements is made, will be considered in this study. Their common oxidation states have electronic configurations ndm where n = 3, 4 or 5 and m = 1-9. For example, copper is transition metal, because Cu(II), 9

3d , is a common oxidation state of copper; whereas zinc is not, as its only common oxidation state, +2, has a filled 3d"*"^ configuration.

Each d-block can be roughly subdivided into (a) early transition elements where complete ionization of all valence electrons may occur [ndm and (n+1) s"*"' ^] , for example, Ti^+, 3d^;

(b) electron-rich transition elements, where ionization of only some of the valence electrons occurs with the rest becoming part of the core. Owing to their reluctance to undergo chemical reactions, and in particular their inertness against mineral acids, the 4d and 5d Group VIII platinum metals and coinage 2

metals are historically also known as "Noble Metals".

The increased nuclear charge in the later elements causes a lowering of the energies of the atomic orbitals, in this case, the d orbitals, and raises the successive Ionization

Potentials. Hence the highest stabilized oxidation state of 7 1 copper is +4 (3d ), as in the anionic fluorocomplex CS2fCuFg]

Some general contributing factors to the stabilization of

the common oxidation states are: (1) Electronic contributions where shells that are empty, half-filled or full are particu-

larly stable, such as Ti , d ; Mn , high spin d ; Ag , d

(2) Thermodynamic factors where maximum Ligand Field Stabili•

zation Energy (LFSE) is obtained. Some examples for octahedral 2+ 8 3+ 6 systems are Ni , d ; Co , low spin d . (3) Kinetic factors

where complexes are rather inert or reluctant to undergo ligand

substitutuions such as the octahedral complexes of Cr(III) and

Co(III) .

The uncommon oxidation states are those without some or

all of the above mentioned stabilization factors, though there

is no sure guide to the relative stabilities of the various

oxidation states,either of a particular element or within a

group. For example, in Group IB, the most common oxidation

state of copper is +2, while +1 is the most common for silver,

and +3 is the most common for gold, for which +1, +2 and +5

are: rather rare.

The unusual oxidation states often lead to many interesting 3

applications. Metals in their higher oxidation states, as in Pt^Fg or Ag*IF2/ are strongly oxidizing and for fluorides,

strongly fluorinating agents; whereas metals in the low oxida• tion states, especially the Group VIII metals, have been found useful in homogeneous and heterogeneous catalysis. It is the higher oxidation states that are of main interest in this study.

The oxidation state of an element is often inferred from

its chemical formula, but the empirical formula may sometimes

be misleading. For example, PdF^ is shown by neutron diffrac•

tion to be a mixed valence complex consisting of Pd^+ and

IV 2- 2 -(Pd Fg) units , and M0CI2 is in reality a complex containing 4+ 3 metal-metal bonded (Mo^Clg) clusters ; on the other hand, 2 +

NiF2 has a rutile structure with individual Ni ions. Clearly,

the empirical formula does not always indicate the coordination

number nor the type of coordination polyhedron.

The higher oxidation states can be generated by two basic

methods, electrolytic oxidation and chemical oxidation.

(1) Electrolytic oxidation: Electrical- current should be

the most powerful oxidizing agent available provided a suitable

medium can be found for the process. In principal, high

oxidation states can be obtained when sufficiently high voltage

is applied, provided that suitable ligands and anions are

available for their stabilization. In addition, since oxidation

takes place at the anode, the metal ion should ideally be present in the form of an anionic complex according to the general equation:

(m-p)- [ ML ]m - [ ML ] n n + pe

This presents some limitations to the use of this method.

(2) Chemical oxidation: The most powerful chemical oxidizer commonly known is elemental fluorine, as illustrated by the abundance of higher valent metal fluorides. However, there are some fluorinated compounds where, the element-fluorine bonds are even weaker than the F—F bond in elemental fluorine, 4 5 such as KrF~ or (Kr_F.,)(AsF,) and PtF_ . These are known z z J b o to be strong oxidizers also capable of complexing the species produced, thus generating the novel heptavalent halo-complexes otherwise unobtainable by simple elemental fluorination ^.

Examples are:

+ (BrFg) (AsF6)~ + 2 Kr + F2 and

+ + - 8,9 2PtF6 + 2C1F5 (ClF6) (PtF6) + (ClF4) (PtFg)

These examples also illustrate the convenience of a chemical oxidation reaction when the oxidizer is incorporated into the product as a fluorido ion. 5

The highest oxidation states of the 4d and 5d transition metals are generally found in the binary fluorides; some exceptions are found in ruthenium and osmium, as RuO^ and OsO^ . The non• existence of the octavalent ruthenium and osmium fluorides is perhaps due to steric limitations of the coordination sphere.

The chemistry of transition metal fluorides has been extensively studied and reviewed over the years . One of the interests in our group is the exploration of the largely undeveloped area of transition metal fluorosulfates, where the transition metal is an electron-rich 4d or 5d element.

The fluorosulfate system was chosen for this study for a number of reasons.

1) The anion has often been termed a pseudohalide because of the resemblance of its chemistry to those of the halides, in 16 particular, that of fluoride . In fact, the fluorosulfate ion satisfies all except one of the requirements suggested by 17

Cotton and Wilkinson for a pseudohalide.

2) It allows the use of a variety of synthetic methods, and hence provides some flexibility. Furthermore, the existance of

a stable and versatile oxidizer, F02SO—OSO2F, makes the syntheses of compounds with the metal in the high oxidation states more promising. This reagent, bisfluorosulfuryl peroxide, is a sym• metrical combination of two fluorosulfate radicals, in good analogy with the halogen molecules.

3) Like the fluoride, the fluorosulfate ion is capable of polydentate coordination, but unlike the fluoride ion, it can be conveniently studied by vibrational spectroscopy in the solid state.

4) Like hydrogen fluoride, fluorosulfuric acid, HSO^F, has been extensively studied as a non-aqueous solvent and as a 18 — 21 synthetic reagent . Fluorosulfuric acid, unlike HF, has a wide liquid range and hence provides a good reaction medium.

5) As the anion of perhaps the strongest simple monobasic 19 acid, HSO-jF , fluorosulfate is highly electronegative and is a weakly coordinating ligand, and hence should be suitable as a counter-ion for the highly electropositive high-valent metal ions .

6) The study can be further extended to similar anions with related oxyacid groups -SO^X, where X = CF^, Cl, etc.

The chemistry of the anion trifluoromethylsulfate, CF^SO^ , 22 has been reviewed in detail very recently

B. THE FLUOROSULFATE RADICAL AND ITS ANION

1. The Fluorosulfate Radical

The existence of the fluorosulfate radical in equilibrium

with its dimer, FO-SO—OS09F, has been suggested from the 7

reversible 'appearance of a brownish yellow coloration when bis(fluorosulfuryl) peroxide is heated. This is a phenomenon

reminiscent of the N2C>4 ^ 2 N02 equilibrium. Some of the chemistry of S2°6F2 does indeed involve the SO-jF* radical.

S F n1957 Dudley and Cady, who first synthesized 2°6 2 -*- ' provide 23

supporting evidence from the spectrophotometric measurements of the temperature dependent absorption of the fluorosulfate

radicals at 474 nm, and also from the temperature dependent

pressure measurements at constant volume between 450 K and 600 K.

The enthalpy of dissociation of S20gF2 to two SO^F- radicals is

found to be 97.6 kj/mol and 92.1 kJ/mol by the two methods

respectively. Since then, extensive experimental and theore•

tical studies have been made and are summarized below. The

enthalpy of dissociation of S20gF2 has also been determined by 24

a kinetic study to be 91.3 kJ/mol

The ESR spectrum of the fluorosulfate radical has been

studied between 4 °C and 180 °C A wide single structureless

component with g = 2.0108 is observed at room temperature,

suggesting a single paramagnetic component. The signal inten•

sity is found to vary directly with the intensity of the yellow

coloration of the sample, both of which are temperature depen•

dent. Analysis of ESR intensity as a function of temperature

gives the enthalpy of thermal dissociation of S„0,Fo a value z b z of 93.8 kJ/mol. 8

King, Santry and Warren have studied the high resolution electronic spectrum of the fluorosulfate radical-dimer system 2 6 in detail . The observed bands have been assigned to tran- 2 2 sitions between the non-degenerate A2 ground state and the 2 and E excited states. The visible absorption centered at 2 2

516.0 nm has been assigned to the A2 —*• E transition. The proposed assignment is supported by their CNDO molecular orbital 27 calculations of the energies of the lower doublet states 2

The assignment of A2 as the ground state is also consistent 28 with another independent calculation . Vibrational analysis 29 of the 516.0 nm absorption system gives the six fundamental vibrational frequencies of the fluorosulfate radical with C 3v 2 2 symmetry. (The assigned fundamental frequencies for the A

ground state are at 1055.5 cm ^, v2 at 839.3 cm at

533.5 cm 1, all of a^ symmetry; and of e symmetry: at

-1 -1 -1 1175.5 cm , v5 at 604.1 cm , and vg at 369.4 cm ). Such an assignment is further supported by a normal coordinate analysis of the fluorosulfate radical using an Urey-Bradley force 30 field . More recently, a detailed matrix isolation study of the fluorosulfate radical from the pyrolysis of S2°6F2 confirms 2 9 the vibrational assignments by King and Warren , as shown

31 in Table 1 . Besides the expected SO-^F' radical, S03, HS03F and S~C» F„ are also observed in the argon matrix at 14 K, 2. 5 2. each identified by their individual matrix isolated spectrum. TABLE 1

3 FUNDAMENTAL FREQUENCIES (cm ) OF MATRIX ISOLATED S03F AND COMPARISON WITH OTHER RESULTS

S03F Ground State ( A2) S03F Excited State ( E) S03F

Fundamental IR Matrix31 IR Matrixb Vapor29 Matrix . Vapor29 Raman5 Argon Nitrogen /C Solution

e v S-0 Stretch 1053.0 l 1390 1055.5 947 966 952. ,9 1082

f v S-F Stretch 2 832.8 815 839.3 796 820 800. ,5 786

v (a^ S-O 531.2 3 Deformation 585 533.5 512 511 515. 0 566 v4 (e) S-0 Stretch 1177.4 1550 1177.5 1114 . 5 1287

v (e) S-0 601.0 5 Deformation 650 604.1 505. 7 592

V (e) S-F Wag 366.0 380 6 369.4 346. 9 4 09

a. reproduced from reference 31. b. K.A. Oakes, Phd. Thesis, University of Washington, 1972. c. The E mode fundamentals of the E excited state could not be clearly identified since transitions to these vibronic levels were much weaker than those to the A, levels. d. Taken of an aqueous solution of the sodium salt. H. Sievert, 2. Anorg. Allg. Chem., 289,15(1957) e. Identified as SO, in Ref. 31.

f. Identified as S-0CF_ in Ref. 31. The photoelectron spectrum of the fluorosulfate radical 2 again confirms a ground state. The first adiabatic ionization potential is found to be 12.85 eV. The observed sharp peak indicates the removal of electron from a non-bonding orbital.

The electron affinity (E ) indicates the feasibility of a attaching an additional electron to any species as defined by

X (g) + e *•> X (g) .

The electron affinity is commonly not directly measurable,and no value for the fluorosulfate radical has been reported. An estimate can be made using the standard Born-Haber Cycle 33 calculation on potassium fluorosulfate, for which all other thermochemical data are either available, or (like the lattice energy) can be estimated. The details of the calculation are shown in Appendix B. A rather high value of 1110 kJ/mol or

11.5 eV is obtained from this calculation,as compared to tabula ted values of 333 kJ/mol (3.45 eV) for fluorine and 348 kJ/mol 34

(3.61 eV) for chlorine . The assumptions made in the calcu• lation, ie. the uncertainties in using the Kapustinskii equatio for a NaCl lattice in calculating the lattice energy of KSO^F, and the assumption that the ionic radius of SO^F is the sum of the S—F distance in KSO-.F and the covalent radius of F, do 11

not appear to have a very significant effect on the resulting value. The major contribution to the high electron affinity comes from the high value of the heat of formation of KSO^F.

(AH° KSO^F = -1155 kJ/mol 35). However, this value has been determined reliably from the calorimetric measurements of the reaction of potassium acetate with fluorosulfuric acid in 35 acetic acid by Richards and Woolf

Further literature scans have resulted in some supporting evidences for the calculated electron affinity of the fluoro• sulfate radical. a) The ionization potential of the SO^F anion should be identical to the electron affinity of the radical . From ESCA 36 — experiments , the first ionization potential of SO^F in

solid KS03F is found to be. vLl eV 1060 kJ/mol) . b) As mentioned earlier, the first adiabatic ionization potential of the fluorosulfate radical is found to be 12.85 eV 32

(124 0 kJ/mol) from its photoelectron spectrum c) The first adiabatic IP for perchloryl fluoride, ClO^F, isoelectronic with SO^F , is found to be 13.10 eV (1264 kJ/mol), 3 6 again from its photoelectron spectrum

As a final point, one may compare the electron attachment process of a fluorine atom to form the fluoride ion to that of a fluorosulfate radical to form the corresponding anion. In the first case an electron is added to a localized atomic 12

orbital, in the other case the receiving molecular orbital 1 is nonbonding and delocalized. Hence, it appears that the electron affinity of the fluorosulfate radical may be much higher than originally expected, and the calculated value of 11.5 eV (1110 kJ/mol) should be considered as a reasonable estimate,though it may be somewhat high. The calculated electron affinity of SO^F* is used subsequently, but with caution.

Pauling's original qualitative definition of electro• negativity is "the power of an atom in a molecule to attract electrons to itself". The relative electronegativities scale serves to illustrate such electron drawing abilities through polar bonds, and can be extended to more complex moieties such as the fluorosulfate group.

The electronegativity of the fluorosulfate group may be 119 estimated from the isomer shifts in the Sn Mossbauer 2- spectrum of octahedrally coordinated (SnX^) complexes. 37 Herber and Cheng have found a linear relationship between the isomer shifts 6 (mm/sec) and the sum of Pauling 3 8 electronegativities, while Clausen and Good have shown that 39 the average Mulliken's electronegativities may also be used. The suitability of using either system has been discussed 40 by Huheey and Watts . It appears that a simple correlation 2- can be made between the isomer shifts of the [SnX,] complexes 13

and the corresponding electronegativities. The isomer shifts

of K2[SnFg] and K^SnClg] are -0.43 mm/sec and +0.48 mm/sec 41 2-

relative to SnC>2 respectively . A number of [Sn(SC>3F)g ] 42 119

complexes have been characterized and the Sn Mossbauer

isomer shifts of K„ [Sn (SC> F) r ] is found to be -0.26 mm/sec, with 2. Jo an uncertainty of ±0.03 mm/sec, generally common to most reported values of 6. Since the Pauling electronegativity {y. ) P for F and Cl are 3.98 and 3.16 respectively, AXp(F—Cl) would 2-

be 0.82. AS for these SnX^ complexes would be 0.91 mm/sec

between F and Cl and 0.17 mm/sec between F and SO^F. From a

linear correlation, AY between F and SO^F should be 0.15, Ap 3

giving a value of 3.8 3 for the Pauling electronegativity of

the fluorosulfate group.

Alternatively, the electronegativity of the fluorosulfate 39

group may be calculated using Mulliken's definition , which

is the average of the ionization potential and the electron

affinity. From the values of IP and E discussed earlier, a

the Mulliken electronegativity would be 1/2(12.85 + 11.50)eV,

which is 12.18 eV. Conversion to the Pauling units is possible X (M) = 0.168 (IP + E - 1.23) 43 via the equation

A value of 3.88 is then obtained for the fluorosulfate group, 14

in good agreement with the "Mossbauer Value". The agreement may be fortuitous, as strictly speaking, the relative numerical values of the electronegativities can only serve as semi• quantitative estimates. Nevertheless, it is clear that the fluorosulfate group is highly electronegative, the value being much closer to fluorine than to chlorine. * 44

The Taft inductive constant (a ) is a measure of the electron inductive effect. For a number of isostructural tin compounds of the type XYSnCSO^F^ with bridging bidentate 119

SO^F groups, the observed Sn Mossbauer quadrupole splittings

A (mm/sec) are linearly dependent upon the sum of the Taft inductive constants of the axial ligands X and Y, where X and Y 4 5 are CH^, Br, Cl, F or SO^F . The Taft constant of the fluoro• sulfate is found to be 3.68 as compared to 3.08 for fluorine,

2.94 for chlorine, and 2.8 0 for bromine. This indicates the fluorosulfate group has a greater ability than F to withdraw electronic charge via a and TT effects and delocalize the charge over the entire group. 2. The Fluorosulfate Anion, SO^F-

Theoretical calculations have suggested that the electronic

1 2 7 28 ground state of the fluorosulfate anion is A^ ' . The vibrational frequencies of the anion determined from a normal coordinate analysis are, for the a^ type: v-^, 1082 cm ,

\>2t 786 cm 1, v^, 566 cm 1; and of the doubly degenerate

1 1 vibrations, v^, v,., and vg are 1287 cm , 592 cm , and

409 cm 1 respectively.

The ligand field parameters of the SO^F measured from the reflectance spectra of nickel salts have been compared to 4 6 those of other weak field ligands . Values of the ligand field splitting parameter (Dq) and the interelectronic repulsion parameter (B) for SO^F , F , Cl and H^O are listed in Table 2.

The values show that the fluorosulfate ion is a weak field ligand, much more like fluoride than chloride, their relative positions in the spectochemical series being Cl~ < F~ ^ SO^F-

< H20. The value of B is reduced from that of the free ion on coordination. The greater reduction of B with SO^F than with the fluoride reflects the better ability of the fluoro• sulfate ion to delocalize electronic charge. Hence the relative positions in the nephelauxetic series would be Cl~

< SO_F~ < Hn0 < F~. 16

TABLE 2

LIGAND FIELD PARAMETERS FOR SOME NICKEL SALTS

Salt S03F~ F~ C1~ H2°

Dq 734 730 720 850

B 905 960 760 940

X-ray crystal structure studies have been made on the 47 48 ionic fluorosulfates, KS03F and NH4S03F . Unlike the fluorides, the SO^F anions are arranged in a disordered fashion in the KSO^F crystal and partially disordered in NH^SO^F. The symmetries of the anions in the orthorhombic crystals appear to

be C„ in KS0oF. In NH.SO,F, 75 % have C- symmetry2 1 , with the 2v 3 4 3 2v

remaining of C^v symmetry.

The fluorosulfate anion is often compared to a number of

similar weakly coordinating anions, like perchlorate, C104 ,

F and difluorophosphate, P°2 2 tetrafluoroborate, BF4 . In general, these anions are of comparable size and often isostructural in simple ionic compounds.

Numerous attempts have been made to study the chemistry of perchlorates in detail, though some were hindered by the low thermal stabilities of some of its compounds, such as the 4 9-51 halogen perchlorates ; the strongly oxidizing nature of 17

its parent acid, anhydrous HCIO^ J , and the dangerous self- dehydration process of anhydrous HCIO^ to give the potentially 52 explosive dichlorine heptoxide, C^O^

Difluorophosphoric acid, the parent acid of PO2F2 , is highly viscous, and relatively difficult to purify. It is also 53 54 a surprisingly weak acid, in fact, a base in r^SO^ ' 55

Tetrafluoroborates have been extensively studied even

though there is no useful anhydrous protonic parent acid. The

fluorosulfate system stands out mainly because of its well known radical-peroxide chemistry and the convenient handling of its parent acid, HSO^F. 3. Summary

As outlined in the preceding sections, the fluorosulfate

radical and anion display a great similarity to fluorine and

the fluoride ion. Some of the above mentioned parameters of

the fluorosulfate group are compared to those of fluoride in

Table 3. In view of their respective transition metal chemistry,

some contrasting points may be singled out.

i) Based on the difference in size, higher-coordination numbers

as well as oxidation states are expected for transition metal

fluorides in binary systems.

ii) Transition metal fluorides should display higher thermal

stabilities in comparable compounds, thus allowing high tempera•

ture synthesis. 18

TABLE 3

SELECTED PARAMETERS FOR THE FLUORIDE

AND THE FLUOROSULFATE SYSTEMS

Anion

S03F

IONIC RADIUS [pm] 136 222 (236)

ELECTRON AFFINITY [ kj/mole] * 333 1110

ELECTRONEGATIVITY [PAULING] 4 .00 3.83

ELECTRONEGATIVITY [MULLIKEN] 1 3.91 3.88

TAFTS INDUCTIVE CONSTANT O* b 3 .08 3 .68

LIGAND FIELD PARAMETERS a

FOR OCTAHEDRAL Ni2+ Dq(cm 1) 730 734

B (cm-1) 960 905

Anion

Radical 19

iii) Transition metal derivatives with the metal in low or intermediate oxidation states may be better achieved in the fluorosulfate system. iv) The larger size of the SO^F groups will make magnetic interaction via the superexchange mechanism more difficult.

Hence fewer magnetically concentrated fluorosulfates are expected. v) The magnetic and spectroscopic properties of a transition metal ion in low or intermediate oxidation states should approach the free ion situation more closely in fluorosulfates than in fluorides. vi) The principal structural tool for fluorosulfates would be vibrational spectroscopy, especially on polycrystalline powders, replacing X-ray diffraction as an important method. vii) While comparable synthetic reagents are available, the fluorosulfate reagents seem less corrosive and permit the use of glass apparatus throughout, while the fluoride reagents require metal or fluorinated plastics.

C. SYNTHETIC ROUTES TO FLUOROSULFATES

The chemistry of fluorosulfates has been fairly extensively 16 56 57 studied and reviewed in recent years ' ' . The general synthetic routes commonly employed that are most suited for the 20

preparation of transition metal fluorosulfates are described below.

1. Direct Insertion of sulfur trioxide into metal fluorides

such as the reaction of SO^ with BaF2 at 200 °C to give barium 5 8 bisfluorosulfate :

200 °C

BaF2 + 2 S03 »- Ba(S03F)2

While this method is one of the earlier methods used ^9,60^ ^ is best used for the preparation of ionic fluorosulfates. Incom• plete insertions into many transition metal fluorides have resulted in mixed fluoride-fluorosulfates. As an example, only 61

4.5 moles of S03 were taken up by one mole of WF^

2. Reactions In Fluorosulfuric Acid

(a) Metal salts such as chlorides and carboxylates ^ ^,

react with HS03F according to the general formula

MX + nHSO-.F — M(SO-.F) + nHX. n 3 3 n

In order to achieve complete conversion, it should be possible to remove HX from the system, either as a volatile byproduct, such as HC1, e.g.

44 SnCl2 + HS03F (excess) »> Sn(SC>3F)2 + 2 HC1 or a soluble protonated species, as in

Cu(CH-.C00)o + HSO..F (excess) Cu(S03F)2 + 2 [CH3C (OH)2 ] (SO

This method is limited by the choice of suitable precursors, which are often available only in the lower oxidation states.

These precursors are often unreactive, in particular where polymeric or layer structures are found, e.g. CrCl^ ^ .

(b) The solutions of pure metals in HSO^F were investigated by Brazier and Woolf Many metals were found to be inert to even boiling acid, while those that reacted, such as Cu,

Bi, Na, K, Ca, In, Tl, gave fluorosulfates in the lower oxidation states only.

45 (CH3)4Sn + 2 HS03F (CH3)2Sn(S03F)2 + 2 CH4

(d) The reactions with transition metal carbonyls should cleave the metal-carbon bonds, resulting in fluorosulfate complexes in which the metal is in a low oxidation state. 3. The Silver Salt Method

This is commonly used to synthesize oxyacid derivatives according to the general scheme:

MXn + n AgS03F w-> M(SC>3F)n + n AgX where X = Cl or Br, and n = 1 or 2.

The formation of the silver halides provides the basic driving force. Limitations are the choice of solvent and the avail• ability of the starting material. The non-oxidative nature of this method has found applications in the syntheses of metal carbonyl fluorosulfates such as

M(CO)5X + AgS03F M(CO)5(SC>3F) + AgX

M = Re, X = Cl 6V; or M = Mn, X = Br 68,69

4. Reactions with Oxidizing Agents

These involve the reactions of bis fluorosul furyl peroxide,

F 3rom; ne ^2^6 2 ' ' '- fluorosulfate, BrS03F or chlorine

56 fluorosulfate, C1S03F with a variety of substrates such as chlorides, oxides, or even the elements themselves. Some examples are: 23

120 °C 70 SnCl4 + S206F2 Sn(S03F)4 + 2 Cl2

R.T. A _

Au + BrS03F (ex.) »- Au (S03F) 3 • 2BrSC>3F Au(S03F)3

This route appears to be the most promising,when attempting the syntheses of compounds with the metal in a higher oxida• tion state.

The two reagents which have been used most often have

rather contrasting chemical properties. While S20gF2 is the

better oxidizer, BrSC>3F is an excellent solvolyzing agent. A

comparison with F2 and BrF3 as their fluorine analogues seems fair. Hence passivation, resulting in incomplete or rather

slow reactions is often encountered when using So0,Fo, which Z D Z interestingly is synthesized in a copper reactor at ^ 180 °C

(Section II.C.2). On the other hand, excess BrS03F used as solvent is often difficult to remove, and may involve side reactions with the halogens produced or further complexation 71

as in the Au—BrS03F system described by Johnson, et al.

Such a system has been reinvestigated very recently in our laboratory, resulting in the identification of the initial III 7 2 complex as Br^[Au (S03F)41 . Some physical properties of HSO-.F, S„OcF„, and BrSO^F are listed in Table 4. 3 2 6 2 3 24

TABLE 4

SOME PHYSICAL PROPERTIES OF HS03F, S20gF2 AND BrS03F

19 56 57 Compound HSO-.F So0^F_ BrSO-,F 3 2 6 2 3

boiling point ( C) + 162.7 + 67.1 + 117.3

melting point ( C) - 88.98 - 55.4 + 31.5

density (gm/cm ) 1.726 @ 1.645 @ 2.238 @ 25°C 35.5°C 25°C viscosity (centipoise) 1.56 @

25°C

-4 -4 specific conductance 1.08 x 10 7.21 x 10

(Q, cm ) @ 25°C @ 25°C

Vapor pressure log P =5.49916 log P =8.544 equation mm ^ mm 1.2925 x 10' 2.195 x 10'

T(K) T(K) 25

D. VIBRATIONAL STUDIES OF THE FLUOROSULFATE GROUP

Vibrational spectroscopy has been widely used as a diagnostic tool in the study of coordination modes of fluoro• sulfates. Compounds containing fluorosulfates may be classified as: 1. ionic compounds where SO^F ions are present, 2. covalent monodentate, 3. bidentate, 4. triden• tate, and 5. tetradentate coordinated. Each of the different coordination modes has its own distinct features in the vibra• tional spectra, these are listed in Table 5.

Bonding will occur through the more basic oxygens except

in the tetradentate mode,where fluorine will be involved.

The local symmetry of the fluorosulfate group would be reduced from C,„ to C in the mono- and bi-dentate cases, thus giving jv s nine fundamentals instead of six for C^ symmetry. All fundamentals are both IR and Raman active. The sulfur-oxygen and sulfur-fluorine stretchings are most useful in identifying the coordination mode of the fluorosulfate group, as they show wide positional differences. A correlation diagram of the frequencies and assignments of the different vibrations of the fluorosulfate group is shown in Figure 1.

For ionic fluorosulfates, where all three oxygen-sulfur 47 bonds are equivalent, such as KS03F , six distinct vibrations 6 2 should be observable in both infrared and Raman . However 26

TABLE 5

STRUCTURAL BONDING DIFFERENTIATION FOR THE SO,F GROUP

number

local S03F description of diagnostic vibrational group sym- of bonding funda- mode and frequency metry pattern mentals

ionic 6 v S-F 800-700 cm 1

1 COT7 tridentate 6 v S-F 8 50 cm •3V tetradentate 6 v S—F 65 0 cm

-1 C covalent monodentate 9 v S-0 900-750 cm -1 covalent bidentate 9 v S-03(3rd)1080-950 cm FIGURE 1 Correlation Diagram for Fluorosulfate Group

S v S0 SO 3 vSF 6 s0 Ionic asym 3 sym asym 3 *sym °3 Prock

v3At vbE a) unperturbed V2A! V5E

-1 -1 1 KSO3F*7 1285cm 1084cm 741cm 587cm- 1 571cm" A05cu»" b) perturbed

-1 -1 1 1 1 73 j1 1 -1 -1 570cm" 416cm" 403cm" KOS03F 1278C« 1246cm" 1077cm 755cm 600cm 588cm

1 (1262cm"1) (594cm A) (4090a" )

g S0 S SF YtwlstS02F Ttor6ionS02F 2 8ym 2 v SF Covalent (C,) 'asym'iS0 v S0 v SO bend 2 Yrock °2 rwag

vi, v9A" v6A' Monodentate v7A" viA' v2A* v3A' A' v8A" v5A'

-1 1 1 1 75 1 1 1 1 500cm- 1 390cm" 395cm" F0S02F 1502cm" 1250cm" 788cm" 857cm 577cm" 530cm"

SO3F 6 SO3F 6 SO3F Yrock.S02 TtorsionSQ3F Bridging (C,) v SO3 v S03 v S03 v SF 6

3 Vi»A' v8A" v9A" 6 Bidentate v7A" ViA' v2A' v A' v A'

1 -1 -i 45 1 1 590cm- 1 548cm'• 1 430cm -1 280cm F2Sn(S03F)2 1420cm" 1101cm" 1068cm" 855cm 630cm

3 S03 v SF SO SO 3 Prock Trldencate (C3V) asym vSvm S03 asym 'sym

v3Ax vfcE Bridging V..E v2Ai vsE

1 -i ll 78 -1 -1 568cm 420cm -1 C0 (S03F)2 1265cm 1109cm" 850cm 610cm

Tetradentate (C3V) S0 v SO, v SF fiasym S03 6sym S03 Prock *asym 3 sym

V5E v3Aj Bridging VJAJ v2Aj v6E

79 -1 1 -1 -1 660cm 592cm"1 579cm" 390cm- 1 T13C110(S03F)2 1248cm 1082cm 28

perturbations are often observed in the ionic systems which are caused by (a) non- spherical cations such as N0+ in 73 NOSO-jF as shown in Figure 1; or (b) polarizing cations such + , 74 - as Ag or Li"*" ; or (c) site symmetry effects where the SO^F

ion is in a crystallographic site of lower symmetry than C^v as in the free SO^F ion. These cause splittings of the doubly degenerate E modes. If however the unit cell contains more than one type of fluorosulfate ions with slightly different local

symmetries due to the orientations of cations to anions, all or most of the fundamentals would be duplicated, regardless 64

of symmetry, as in Sr(S03F)2

For a covalent monodentate fluorosulfate group, nine

fundamentals are expected due to the splittings of the three

E modes to pairs of A1 and A" modes when the symmetry changes

from C^v to Cg. The bonding of an oxygen to another atom X

causes the vS—O of S—0—X to shift to lower energies, as in

75 -1 -1 the case of FOS02F , from 1084 cm of KSC>3F to 788 cm ; while the average value of v and v SO- of the unbonded 3 asy sym 2

oxygens shifts to higher values. Furthermore, the electron

drawing effect of the highly electronegative fluorine atom

bonded to sulfur causes the S—F stretching to increase from

-1 -1 741 cm of KS03F to 857 cm as well.

If the fluorosulfate group is covalently bonded in a biden- 45

tata. fashion, as in F2Sn(S03F)2 , the number of fundamentals 29

remains nine because of the unchanged Cg symmetry, but this case can be distinguished from the monodentate group by the in• crease in frequency for the lowest vS-0 vibration. For example, 7 6

the lowest S—0 stretching of (CH-^)2 Sn (SO^F)2 appears at

1072 cm 1, while the two higher S—O stretchings also appear at lower frequencies, 1350 cm 1 and 1180 cm 1. The existance

of bidentate bridging fluorosulfate groups in (CH^) 2Sn (SO-^F) 2 77 has been confirmed by a crystal structure study

In general, the three S—O stretching modes of a monoden•

1 tate group appear between 1500 and 1400 cm for SC>2 asymmetric,

-1 -1 1250 and 1200 cm for SC>2 symmetric, and 950 to 780 cm for

S—O of S—0—X. For the bidentate fluorosulfate group, the three modes are found between 1400 and 1300, 1180 and 1100, 1080 and

950 cm 1. The S—F stretch in both cases would be approximately at

850 cm 1, much higher than that of the ionic fluorosulfate.

This is mainly due to the increase in pir-dIT back donation commonly observed on coordination.

The increase in S—F stretching frequency also holds for the tridentate case and is in fact the major diagnostic sign in distinguishing it from the ionic mode, whose S—F stretching appears in the 750 cm 1 region. Otherwise tridentate fluoro- II 78

sulfates, such as Co (SO^F)2 , would have C3v symmetry giving rise to six fundamentals and S—O stretching frequencies similar to the ionic ones. At present, the only known case of a tetradentate bridging

7 9 fluorosulfate group is the compound Ti^Cl^Q(SO^F)^ . The v v asy SO-,3 and sym SO-,3 frequencie^ s are agai^ n similar to the ionic and tridentate fluorosulfates, but unlike these two types, the S—F stretching frequency appears at 660 cm 1. This remarkable shift is expected, because the S—F bond is weakened as a result of the fluorine-titanium bonding.

In actual experimental measurements, fluorosulfates are generally excellent Raman scatterers, except where the dark colour of the compound or the occurrence of fluorescence present experimental obstacles. The other experimental dif• ficulty often encountered is attack on mulling agents and IR window materials due to the high reactivity of the fluorosul• fates .

E. TRANSITION METAL FLUOROSULFATES

Compared to the fluorides, only limited work on binary fluorosulfates of the transition metals has been done '

The existing binary fluorides and fluorosulfates are shown in

Figure 2. There is an obvious lack of corresponding fluorosul• fates of the second and third transition series,while the majority of the work performed concentrates on the 3d elements. MB IV B VB VIB VIIB -VIII— IB IIB

3 3,4 3,4,5 2,3,4,5 2,3 2,3 2,3 2 2 2 Sc Ti V Cr Mn Fe Co Ni Cu Zn © © ©© © © © © 2 3 4 3,4,5 3,4,5,6 5,6 3,4,5,6 3,4,5,6 2,3*4 1.2 Y Zr Nb Mo Tc Ru Rh Pd Ag Cd ©

3*4,5,6 3,5 1.2 3 4 3,5 4,5,6 4,5,6,7 4,5,6,7 3,4,5,6 La Hf Ta W Re Os Ir Pt Au Hg © © ©

* denotes a mixed oxidation state + 2, + 4

Oxidation states of existing binary transition (Fiuorides~uPPer rowS) . , n ir • (Fluorosulfates—-in circles) metal fluorides and fluorosulfates 32

Therefore it appeared appropriate to investigate the later

transition elements in this study. Previous reported work is now summarized as follows.

The displacement reactions in fluorosulfuric acid of 6 3

several transition metals salts has been investigated . . The anhydrous bisfluorosulfates of Mn, Fe, Co, Ni, Cu, Zn, Cd have been prepared from the acetates by this method. The ease of

2 displacement is in the order CH3C02~ > SO^ ' > Cl~ > F~.

Using the anhydrous chloride as starting material, Goubeau and Milne prepared the corresponding Cu(II), Zn(II) and Fe(III) 6 2

fluorosulfates from fluorosulfuric acid . Potassium chloride, which is solvolyzed to its fluorosulfate, must be added to

aid the removal of these acid insoluble fluorosulfates from

the metal chloride surfaces. Although no crystal structure

has been reported on these compounds, detailed magnetic and

spectral (vibrational and electronic) studies have been made on the divalent fluorosulfates 7^,64^ These compounds are

expected to have tridentate bridging fluorosulfate groups around

the octahedrally coordinated metal centers. In the Cu(II) and

Mn(II) compounds, the fluorosulfate groups probably have two

stronger and one weaker metal-oxygen interaction, which results

in lowering of the symmetry of the group.

Silver(I) fluorosulfate was originally synthesized by the

reaction of silver metal with nitrosyl pyrosulfate, (N0)9S90_, 33

8 0 in bromine trifluoride . However, the solvolysis of silver trifluoroacetate, AgOCOCF^, in HSO^F has been found to be a

81 66 much more convenient synthetic route . Brazier and Woolf have suggested the possible existence of silver(I) fluorosulfate when the metal is dissolved in boiling HSO^F, although the

product so obtained contains large amount of Ag2S04 as impurity. 8 2

The vibrational spectrum of AgSO^F suggests ionic fluoro• sulfate groups with some distortion due to the high polarizing power of the silver cation.

The reactions of gold and platinum with bromine(I)

fluorosulfate, BrSC>3F, give Au(S03F)3 and Pt(S03F)4 through some intermediate complex formations in the reaction medivuns-

These complexes can be thermally decomposed to yield the pure products. Oxidative reactions with ^2°6F2 nave yielded tran• sition metal oxyfluorosulfates. Metal oxidations have produced

83 84 Re03(S03F), Re02(S03F)3 , and Mo02(S03F) . Reactions ,

with carbonates or oxides have resulted in MnO(S03F), CoO(S03F), 8 5

NiO(S03F), and Ag20(S03F)2 . From the reactions with their O C hexacarbonyls, Mo02(S03F)2 and WO(S03F)4 are obtained. When

S F S reacte 2°6 2 "*" ^ with pentachlorides of Group VB, M0(S03F)3 8 3 with M = V, Nb, or Ta , are obtained. The conversion of

87 Cr02Cl2 into Cr02(S03F)2 by S206F2 is also reported . 8 8

Just before the initiation of this study, Brown and Gard published a short note on the synthesis of a chromium(III) fluorosulfate from the reaction of sulfur trioxide and chromium pentaf luoride . Interestingly, bis fluorosul furyl peroxide

S2°6F2' was f°und to ke a by-product in the above reaction. 8 9

While this work was in progress, Brown and Gard further pursued the synthetic route of oxidation of metal carbonyls with S^O^F^. They have reported the syntheses of some metal fluorosulfates, oxyfluorosulfates and carbonyl-fluorosulfates

formulated as Mn(SC>3F)4, VO(S03F)3, FeO(SC>3F) and Fe (CO) 4 (SC>3F)

starting from Mn„ (CO), n, V(CO), and Fe(CO)c-.

E. STRUCTURAL CHARACTERIZATIONS OF TRANSITION METAL

FLUOROSULFATES

The most powerful means of structural identification would be single crystal X-ray diffraction analyses. However, most fluorosulfate compounds do not form suitable crystals, hence spectroscopic methods are commonly employed for structural characterizations.

Among them is the already discussed vibrational (infrared and Raman) spectroscopy. In addition, the magnetic properties of the complexes are studied by bulk magnetic susceptibility 90 91 92 measurements ' and electron spin resonance spectroscopy 93-95

The electronic structures of the metals are studied by solution visible-ultraviolet, solid state mull, and diffuse reflectance spectroscopy. 35

II . EXPERIMENTAL

A. APPARATUS

Since the products and most of the reagents used were hygroscopic, special precautions were taken to avoid contact with moist air at all stages of experimentation. Hence

gases and volatile liquids were manipulated on a vacuum line, whereas solids and non-volatile liquids were handled in a dry box. In addition, chemicals such as chlorine dioxide

and many silver(I) compounds were light sensitive to some

extent. Care was therefore taken to avoid such chemicals

from prolonged exposure to light. A fumehood with darkened

windows, deeply tinted glass containers or aluminum foil

wrappings were employed in such cases.

1. Pyrex Vacuum Line

A general purpose pyrex vacuum line,about 60 cm long with

five outlets fitted with Kontes Teflon stem stopcocks, was used.

A standard taper B19 ground glass socket at one end connected

the line to a mechanical rotary vacuum pump (Welch Duo-Seal

Model 1405) via a pyrex connecting piece with a safety trap

cooled with liquid nitrogen, to protect the pump from corrosive 36

volatile materials. All other outlets of the line were BIO ground glass sockets, spaced evenly along the manifold. The connecting piece had two other outlets between the safety trap and the manifold: one connected to a mercury manometer via a stopcock and BIO ground glass connection; the other to a stopcock and BIO socket often used to pump off unwanted volatile reaction materials. Another outlet between the safety trap and the vacuum pump served as a leak-valve to the atmosphere via a Kontes glass vacuum stopcock. When transfer- ing liquids with lower volatilities, it was found convenient to use a T-connecting bridge rather the entire manifold.

Such a T-piece consists of simply BIO sockets at either end and a BIO cone connecting to the main manifold via a Kontes

Teflon stem stopcock. Typical vacuum generated on such a line -2 is of the order of 1 pascal (> 10 torr).

2. Metal Vacuum Line

In the reactions where elemental fluorine is involved, a metal vacuum line was used. This was constructed with

1/4 inch O.D. monel tubings equipped with Whitey valves(1KS4316) and was operated in the manner similar to the pyrex line.

Copper tubing (1/4 inch O.D.) was used for connections to

the manifold for more flexibility. The pressure in the

system was monitored by a Crosby pressure gauge and a NRC 801 37

thermocouple vacuum gauge. The basic handling and manipulation techniques on the line were similar to those described by 9 6

Peacock . Other specific metal lines used for the preparation

of S20gF2 and CsAgF^ are described in Section II.C.2.

3. Reaction Vessels

Several types of pyrex reactors were used for most • reactions except where glass attacks were expected. In these cases, either monel or Kel-F reactors were used.

(a) Pyrex Reactors (Fig. 3)

Most often used was a test-tube type reactor of about 25 ml in volume with a constriction leading to a Kontes Teflon stem

stopcock, and a side arm extending to a BIO ground glass cone

(Fig. 3a). Larger vessels of this type up to ^500 ml were also

used for storage of volatile reagents. Where high reaction

temperatures (up to 150°C) were employed, and volatile liquids

or large amount of qaseous products were expected, 2 mm thick

wall variations of above were used. These could stand up to

internal pressure of 7-8 atmospheres (Fig. 3b). The Teflon

stem stopcock served as a pressure releasing safety valve as

well. Another variation to enlarge the volume of the container

was fitting round bottom flask to the end of the tubes (Fig. 3c)

These one piece containers can be loaded through the 38

FIGURE 3

PYREX REACTION VESSELS

• TEFLON STEM STOPCOCK Q- 8-10 GROUND GLASS COU£

PYREX GLASS

a) one-part reactor and b) one-part thick-walled storage vessel reactor FIGURE 3

PYREX REACTION VESSELS

TEFLON STEM STOPCOCK

B-lO GROUND GLASS CONE

Two-part reactor c) one-part reaction bulb 40

constrictions with the aid of a small diameter funnel. They

have the advantage of excluding stopcock grease and withstand•

ing a positive internal pressure. However, solid products

that did not come loose on shaking or tapping, could not be

removed without breaking the vessel inside the drybox. Further•

more, "bumping" of a suspension or solution often took place while removing the volatile liquid from the solid product by

vacuum distillation. Thus, proper vacuum seals were impeded

by solid stuck between the Teflon stems and the glass surfaces;

while loss of product from the vessel resulted in losing track

of total product yield, where the progress of the reaction was

monitored by weight. The two part pyrex reactor used consists

of a 50 ml round bottom flask with a B19 ground glass cone and

a corresponding adaptor top with a Teflon stem stopcock between

a B19 socket and a BIO cone (Fig. 3d). This type of reactor

had the obvious advantage of easy loading and removal of solid

material but was not suited for reactions at high temperature

or high pressure, and stopcock grease contamination was a

problem as well. However, problems arose from "bumping" could

be avoided by having a sintered glass filtering disc between

the B19 joint and the Teflon stopcock of the upper adapter.

Teflon coated magnetic stirring bars were used in all these

reactors to facilitate proper mixing of reaction mixtures. 41

(b) Metal Reactor (Fig. 4)

These were ^100 ml two part monel reactors, each held

together with six bolts to withstand high pressures. They were selected for their inertness toward corrosive reagents

(e.g. F2, HF or halogen fluorides). The top part was fitted with a Hoke valve (#4 31) and a swagelock 1/4 inch connector or a standard tapered BIO cone to attach directly to the metal or pyrex vacuum line. A vacuum tight seal was obtained with a Teflon O-ring inserted into a groove between the pot and the

lid. Solid samples can be easily loaded and removed in the

drybox, however the physical changes of the reactants during

the reactions could not be observed in these reactors. A monel

fluorine flow reactor was also used for the preparation of

CsAgF^, which is described in Section II.C.2.

Where high internal pressure and high reaction temperatures

were not involved, but corrosive reagents or reaction products

were encountered, Kel-F reactors were used. They were semi-

transparent and detachable. These consisted of upper monel

fittings similar to the monel reactor; and bottom monel fittings

to retain the Kel-F tube as shown in the diagram. These Kel-F

tubes sometimes had the problem of developing longitudinal cracks

after repeated sudden changes in temperature. Hoke Valve ( No 431)

Monel Metal Tube

Bolts to Secure Lid to Bottom Vessel \ Lid n n ~N n Condenser Inlet feottom

Condenser Inlet

Monel Metal Reaction Vessel (150 ml )

Monel Metal 2-Part Reaction Vessel ( Front View ) Fig. 5 Kel-F Reaction trap 44

4. Fluorosulfuric Acid Distillation Apparatus

Since the HSO^F available commercially was only technical grade, it was essential to purify it by a double distillation 97 process described by Barr et. al. . The distillation appara• tus is shown in Fig. 6. The system was first flushed with

P2°5 dried nitrogen gas and the distillation was carried out at atmospheric pressure in a fumehood. Most of the HF impurity was removed in the first distillation by a counter stream of

dry N2 through the outlet guarded by an anhydrous CaSO^ drying tube. The second distillation yielded a constant boiling fraction,

(163 °C) which was collected directly into an evacuated reactor or A storage container (Fig. 3c) .

5. Trifluoromethylsulfuric Acid Distillation Apparatus

HSO^CF^ was purified by distillation under reduced pressure as the purity of the commercially available acid was not known.

A conventional distillation set-up with a Vigreux column and fraction separator was used. After flushing out the system

with dry N2, the acid was distilled at ^15 torr from concentrated

H2S04. The constant boiling fractions of ^110°C were collected directly into evacuated storage containers, and further trap- to-trap vacuum distilled before use. Fig. 6 Fluorosulphuric Acid Distillation Apparatus. Ul 6. Drybox

The drybox provided a moisture free environment for handling hygroscopic solids and non-volatile liquids. This was a Vacuum Atmospheres Corporation "Dri-Lab" Model HE-4 3-2

filled with K-grade N2- The dryness of the nitrogen was maintained by constantly circulating over molecular sieves.

The molecular sieves were regenerated periodically by a buil in heating unit (VAC. Model HE-93-B "Dri-Train") to ensure maximum dryness. A dish of ^2^5 ^eP^ inside the box also served as a moisture indicator and drying agent.

7. Miscellaneous

(a) Stopcock grease

To ensure air-tightness, all ground glass connections were lubricated with a low volatility grease that must be inert to halogen containing compounds. A halocarbon grease,

Fluorolube GR-90, CF-Cl(CF--CFC1) CF-Cl, from Hooker z z n z Chemicals, was found to be satisfactory.

(b) Balances

As the progress of most of the reactions were followed by mass measurements, several balances were used: Mettler

Gram-atic analytical balance #1-911, with precision and read 47

ability of 0.5 mg to maximum 100 gm load; Mettler Gram-atic

B-5, with precision and readability of 0.1 mg to maximum 200 gm

load; located in the drybox was a top loading Mettler PI60

balance with a precision and readability of 1 mg and maximum

160 gm load; for heavy containers, a triple beam balance

capable of measuring a maximum of 26 00 gm was used.

B. INSTRUMENTAL METHODS

1. Infrared Spectroscopy

Three infrared grating spectrophotometers were used.

Most often used was a Perkin-Elmer 457 grating spectrometer

with a spectral range of 4000 - 250 cm 1. A Pye-Unicam SP1100

grating spectrometer with a range of 4000 - 400 cm 1 was also

used to record spectra at room temperature. All spectra were

calibrated with a polystyrene film. Due to the reactive

nature of the samples, mulling agents and window plates were

often attacked. Hence, depending on the reactivity of the

samples, neat solid powder film between rather inert BaF2, and

less inert AgCl, AgBr, KRS-5 (thallium bromide-iodide) window 48'

TABLE 6

COMMERCIALLY AVAILABLE TYPES OF APPARATUS

Apparatus Manufacturer or Supplier

Welsh Dou-Seal Vacuum Welsh Scientific Company; Pump, Model 14 00 Skokie, Illinois.

Kontes High vacuum glass and Kontes; Franklin Park, Illinois. Teflon-stem stopcocks

Metal high pressure Whitey; Columbia Valve and Fitting and vacuum valves CO., Vancouver, British Columbia. Hoke; Hoke Inc., Creskill, New Jersey. Autoclave; Autoclave Engineering Inc., Erie, Pennsylvania.

Fluorolube Grease, GR-90, Hooker Chemical Corporation; GR-362, Fluorolube Oil, Mo-10 North Vancouver, British Columbia

Tygon tubing Fisher Scientific Co., Teflon coated stirring bars Vancouver, British Columbia.

Linde chromatograph-grade Union Carbide; 5A molecular sieves Redondo Beach, California.

Dri-Lab (VAC) Model Vacuum Atmospheres Corporation HE-43-2 North Hollywood, California. Dri-Train (VAC) Model HE-93-B

Mettler Gram-atic #1-911 Fisher Scientific Co., and P160 Balances Vancouver, British Columbia,

IR window materials Harshaw Chemical Company; (KRS-5,, BaF„, Csl, Cleveland, Ohio. AgBr and AgCl)

Quartz Optical Cells Thermal Syndicate Ltd.; Wallsend, Northumberland, U.K. 49

plates with spectral cutoffs at ^800 cm J', 400 cm J", 300 cm "•" and 250 cm 1 respectively were used. In particular, the

silver halides plates were unsuited for the silver compounds

as redox and exchange coupled reactions often took place.

Nujol mulls were also used where possible. Since most IR

samples were extremely hygroscopic, solid powders were protected

from moisture by wrapping black electrical tape around the edges

of the IR window plates inside the drybox. The IR spectra were

then recorded immediately after taking the sample into the

wet atmosphere. Spectra of gaseous material were recorded

using a Monel cell of 7-cm path length, fitted with AgCl windows

and a Whitey valve.

Infrared spectra at liquid nitrogen temperature were

obtained on a Perkin Elmer 225 grating spectrophotometer with

a spectral range of 4 000 - 200 cm 1. The low temperature cell 98 99

was similar to that of Wagner and Hornig ' . It consisted

of an evacuable pyrex body fitted with two Csl end windows,

and equipped with a central dewar column with a brass block

at the end. The brass block was constructed to hold a Csl

window at the center of the cell. The solid sample powder was

deposited onto this cold central Csl window using a "spray-on"

technique with nitrogen as the carrier gas. Numerous

trials were often needed to deposit the optimum amount of sample

on the cold window without having the powder falling off or being

blown-off. 50

2. Raman Spectroscopy

Raman Spectra were recorded on a Spex Ramalog 5 spectro• meter equipped with a Spectra Physics 164 argon ion laser. The green line at 514.5 nm was used for excitation. The solid samples were packed into melting point capillaries and sealed with fluorolube grease in the drybox, permanent flame-seals were then made soon as possible.

In addition, a Cary 81 spectrometer equipped with a Spectra

Physics 125 He-Ne laser source using the exciting line at 632.8 nm was used as well. Sample containers used here were 5 mm O.D. flat end pyrex tubes. The power output measuring at the sample (15 to 20 mW) was rather poor, and the scattering efficiency of the He-Ne laser is very low, but some coloured samples gave rather good Raman spectra on the older instrument.

3. Visible and Ultraviolet Spectroscopy

Electronic spectra were recorded on either a Cary 14 or a Perkin Elmer Model 124 spectrophotometer. Solution spectra were taken with spectrosil optical cells of 1 mm and 10 mm path lengths stoppered with Teflon plugs and wrapped with Teflon tapes. Solid mull spectra were obtained using Fluorolube oil mulls between silica windows. Light scattering was compensated for by placing a nujol-soaked filter paper in the reference beam 1^1, 51

Diffuse reflectance spectra were obtained on a Bausch and

Lomb spectronic 600 spectrometer equipped with a visible reflectance attachment and a Sargent SR recorder. Spectra were

measured between 350 nm and 740 nm with magnesium carbonate

as the reflectance standard.

4. Magnetochemistry

The magnetic susceptibilities of the compounds were deter•

mined at a constant field strength of approximately 8000 gauss, 102

using a Gouy apparatus described by Clark and O'Brien

Measurements were made over the temperature range 300 K to 77 K.

The solid powdered samples were packed into flat end 4.9 mm O.D.

pyrex tubes to a height of 9.0 cm, then either sealed with

fluorolube grease,or in the case of the second tube, capped with

an air-tight Teflon cap to prevent hydrolysis. Samples were

tested for field dependence of the magnetic moment at a field

strength of approximately 4 500 gauss. Results obtained from the

measurements at the constant field strength of approximately

8000 gauss was reported strictly on the basis that the change in

weight of the sample in and out of the magnetic field is greater

and hence can be more accurately measured. The sample tubes

and apparatus were calibrated using HgCo(CNS)4 as standard

The effective magnetic moments of the metal ion, ye£f/ was 52

calculated using the equation:

r 2 yeff = 2.828(xS° T)

cor

was where xM the molar susceptibility corrected for diamagnetism, and T was the absolute temperature. The detail calculation to obtain X^°r is shown in Appendix D.

Diamagnetic corrections were obtained from the litera• ture 90,104^ values used were as followed: (in cm3mol 1)

6 6 Ag2+, -24 X lO" Ru3+, -23 X lO"

6 3+ 6 Ru4+ , -19 X lO" Os , -35 X 10"

6 + 6 K+ ' -13 X lO" Cs , -41 X lO"

6 + 6 Sn4+ , -16 X lO" Ag , -24 X lO"

6 6 Pt4 + , -28 X lO" ; bipy, -105 X lO"

S03CF3 -46 x 10

The diamagnetic correction for SO-^F1 was assumed to be identical '3

2- 6 3 X to that of the isoelectronic SO. (40.1 x 10~ cm mol ) . In order to determine whether the complexes obey the Curie-Weiss Law,

r yeff = 2.828 [X^° (T - 0)]

C03T the magnetic susceptibilities data in l/xM M versus T were analyzed by a least-squares method to obtain the Weiss constant,

5. Electron Spin Resonance Spectroscopy

A Varian Associates Model E-3 Spectrometer equipped with a 100 KHz field modulation was used to record spectra at room temperature and at liquid nitrogen temperature. The X-band microwave frequency was calibrated using a Hewlett-Packard 5245L

Electronic Counter with a 525.6nm frequency converter 8-18 GHz.

Powdered solids or solutions were contained in 4 mm O.D. quartz tubes and evacuated before measurements. Powdered solids were also held in melting point capillaries that were tested

to show no spurious signals.

6. Mossbauer Spectroscopy

The Mossbauer spectrometer consisted of a TMC Model 305

velocity tranducer driven at constant acceleration by a TMC

Model 306 wave form generator and phase-locked to a 400 channel

analyzer. Measurements were made at 80 K with the Ba^^SnO^

source at 298 K. Isomer shifts were reported relative to an 119

Sn enriched Sn02 absorber at 80 K. Samples were contained

in brass cells with mylar windows. The confident limit was

estimated to be ± 0.0 3 mm/sec for both isomer shift and quadru-

pole splitting. The Mossbauer spectrum of AgSn(SO.jF)g was recorded by Drs. T.B. Tsin and J.R. Sams.

7. Melting Points

Melting points or decomposition temperatures were deter• mined using a Thomas Hoover capillary melting point apparatus in which both the sample in a sealed glass capillary and the thermometer bulb are heated in an oil bath. The melting points are reported uncorrected.

C. CHEMICALS

1. Commercial Sources

All chemicals obtained from commercial sources and subse• quently used without further purification were of reagent grade or of the highest purity obtainable unless otherwise noted.

These reagents, along with their suppliers and purities are listed in Table 7.

Additional purifications were needed for the others.

Technical grade fluorosulfuric acid was obtained from Allied

Chemicals and doubly distilled as described in Section II.A.4.

Trifluoromethylsulfuric acid, marketed as FluorOchemical 55 TABLE 7

CHEMICALS

Source Chemical Purity

Ventron Alfa Corp, silver -100 mesh, 99.999% ruthenium -80 mesh osmium -60 mesh platinum -60 mesh

Ag2o 99% AgO 94% AgF not given

AgF2 98%

AgN03 99.7%

Ag02CCF3 99%

Ag0S02CF3 99%

Ag2S04 99.9%

"AgS03F" not given

AgBF4 analytical

RuCl3 anhydrous not given RuO„ 99.9%

Ventron Alfa Corp. OsCl, not given CsNO. 99.9%

Mallinckrodt Inc. NaF analytical reagent KI analytical reagent KC1 analytical reagent Ag2C03 analytical reagent P2°5 analytical reagent

British Drug Houses, Ltd. KC103 analytical reagent

H2C2°4*2H2° analytical reagent Tin -20 mesh, 99.97%

J.T. Baker Co. CsCl 99.9%

Fisher Chemicals a , a ' -bipyridine reagent

Allied Chemicals H2S04 96% A.C.S. reagent

Canadian Liquid Air N, dry K-grade

Willow Brook Lab., Inc (CF3S02)20 98% 56

acid FC-24 by Minnesota Mining and Manufacturing Company, was distilled under partial pressure (Section II.A.5). Stabilized

sulfur trioxide, obtained as "Sulfan" from Allied Chemicals, was used without purification in the synthesis of S„0,F_. Z b Z

However, the sulfur trioxide used in other reactions was distil•

led from 6 5% oleum under a N2 atmosphere and subsequently vacuum distilled. Antimony pentafluoride, obtained from Ozark

Mahoning Co., was purified by first purging the material with

dry N2(g) to remove most of the HF, then distilled under N2 and

finally intermittent vacuum pumping until evolution of gas

bubbles ceased. Both technical grade (Air-Products Ltd.) and

98% pure (Matheson of Canada Ltd.) fluorine were passed

through a NaF drying tower to remove the HF present. The NaF

tower was equipped with a heating system to allow regeneration

after use. Analytical reagent bromine obtained from Mallinck-

drodt Inc. was stored over KBr and Po0c to remove Cl„ and H„0, Z D z z

and vacuum distilled before use. Acetonitrile, CH^CN

(Mallinckrodt Analytical reagent) was purified by two successive

refluxing periods of about ten hours each with ?20^, phosphorus

pentoxide, then stored and degassed over Linde 4A molecular

106

sieves . Dichloromethane, CH2C12 (Analytical reagent) was

dried over Linde 4A molecular sieves for several days, degassed,

and vacuum distilled. 57

2. Preparative Reactions

Other starting materials (which were not commercially available) were prepared according to literature methods.

Bisfluorosulfuryl peroxide, S„0,F„, was prepared by the

Z o Z reaction of fluorine and sulfur trioxide at ^180°C in the 107 108

presence of AgF2 according to Cady and Shreeve ' with several modifications. The reaction apparatus is shown in

Fig. 7. A larger reactor of 120 cm in length, a reaction temperature of 180 °C versus 150 °C, and heating the

SO^ to 50°C versus 25°C gave faster flow rates (^150 g/hr versus 6 g/hr) and allowed the synthesis of larger quantities in a relatively short period of time. A relative flow rate of

S03/N2 : F2 in the ratio 2:1 were used. To avoid the conden- 108

sation of potentially explosive by-product FSO^F , the last condensation trap was also cooled by dry ice (-78°C) and not

liquid 02 (-183°C). The first trap was left uncooled to act as an air cooler and to allow observation and removal of

possible non-volatile side-products. The presence of unreacted

sulfur trioxide was indicated by white flaky solids in the otherwise clear and colorless reaction product. It was removed

from the crude product by washing with 96% H2S04 in a separatory

funnel in a fumehood before vacuum distillation. A greenish

product color indicated the presence of FSO-.F and dissolved F». Crosby Pressure Guage NaF Trap * i m II Hi To Flowmeter 9 cm i_J—L_r ^-20cm-*- Copper Glass i i -e- To F cylinder ^1 2

F2 Outlet To Flowmeter

5COml. Pyrex Flask

Reactor Copper

(f) Whitey Valve To Soda- lime Trap B34 B34 B -0- Hoke 413 Valve 34 •Fluorolube Oil Tube Autoclave Engineering Valves

A B C

Fig.7 Apparatus- for the Preparation of S2 Oe F2 59

Both were removed at -78°C by prolonged pumping under vacuum.

Frequent visual checkings on the appearance of the crude product allowed fine adjustment of the flow rates for optimum yields. The copper tubing connecting the outlet of the reactor and the cooling traps were flamed frequently to avoid clogging by condensed material. The final product was vacuum distilled into a Teflon valve storage trap. Its purity was checked by 19 a gas phase IR spectrum and a liquid phase F N.M.R. spectrum.

A commonly occurring impurity, in particular after the product had been washed,was disulfuryldifluoride, S2°5F2' Present ^n

VL% amounts. No attempt was made to remove this material, which was inert in the reactions performed.

Bromine monofluorosulfate, BrSO^F, was prepared by the

S F accor direct reaction of Br2 with a slight excess of 2°6 2 ding 109 to Aubke and Gillespie

Chlorine dioxide was prepared according to the literature method 110 by the reaction of potassium chlorate and sulfuric acid with oxalic acid as a reducing agent. The resulting CO,,

acted as a diluent. The C102 produced was collected at -78°C,

allowing the volatile C02 by-produced to vent in the fumehood.

It was then purified by pumping at -78°C to remove Cl2 and C02, then by trap-to-trap distillation. Because of its explosive

111 nature , the C102 was not allowed to warm up and usually further reacted with S-O,.F_ to form chloryl fluorosulfate, Z b 2 60

112 ClO^SO-F . An excess of So0,F„ was distilled onto purified 2 3 2 6 2 C1C>2 and allowed to warm by itself from -78°C to room tempera• ture overnight. The excess S„0,F_ was then distilled from the

Z D Z much less volatile ClO^SO^F. Silver(I) fluorosulfate, AgSO^F, was conveniently obtained from the solvolysis reaction of silver(I) trifluoroacetate, 81

AgOCOCF.^ in doubly distilled fluorosulfuric acid . After removal of all volatile materials in vacuum, the resulting crude product was washed with a very small amount of HSO^F 6 6 and vacuum filtered. Brazier and Woolf observed that silver metal dissolved in boiling HSO^F to give a colorless solution containing possibly silver(I) fluorosulfate, which could not be isolated without large amount of sulfate impurities.

This appears to be confirmed when silver metal was reacted directly with doubly distilled HSO^F at 70°C for three hours, removal of all volatile materials in vacuum gave a white solid

contaminated with traces of a clear viscous liquid of very low volatility, most likely H^SO^. The white powder was identified

by its IR spectrum to be an impure AgSO^F.

Silver(II) bis(2,2'-bipyridyl)bis(trifluoromethylsulfate), 11 was s n [Ag (bipy)2 1 (^F^SO^^/ Y thesized from its corresponding silver(I) complex by oxidation with AgO in the presence of 113

excess anions according to Thorpe and Kochi

Both potassium and cesium fluorosulfates were prepared

from the corresponding chlorides by reacting with fluorosulfuric acid , or bromine monofluorosulfate. Due to their low thermal stabilities, the alkali metal bisfluorosulfato-

bromate(I) complexes forme<^ ^n {-ne latter route were decomposed by heating the products under a dynamic vacuum, thus giving the corresponding pure alkali metal fluorosulfates.

Cesium tetrafluoroargentate(III), CsAg^^F^ , was prepared by the direct fluorination of stoichiometric mixture of CsNO^ and AgNO^ at ^300°C according to Hoppe . The reaction was carried out in a fluorine flow reaction system shown in

Fig. 8. The fluorine (Matheson) used was passed through a NaF

column to remove any HF impurities before diluting with dry N2-

The entire reactor was taken into the drybox for manipulation of the product.

D. CHEMICAL ANALYSES

1. Elemental Analyses

Carbon, hydrogen and nitrogen analyses were carried out by the analyst, Mr. P. Borda of the Chemistry Department,

University of British Columbia. All other elemental analyses were performed by Analytische Laboratorien (formerly Alfred

Bernhardt), Gummersbach, West Germany. FIGURE 8 METAL FLOW REACTOR FOR FLUORINATION REACTIONS

heating wire Whitey valve

[iwiiiiiiiiiiiiiiiiir iiiiuiittiHWimniii •

L_ F2/N2 9 r 9 monel lousing

copper1 gasket

thermocouple FLUORINE LINE

reactor NaF Trap

N, fluorocarbon f) Oil 2. Oxidation States Determinations

The oxidation state of some divalent silver compounds were determined by a modified iodometric titration method.

Since the reaction of compounds such as Ag(S03F)2 with aqueous acidic potassium iodide solution gave both oxygen and elemental

iodine according to:

Ag2+ + 2 KI »- Agl + | I + 2 K+

2+ + + and 2 Ag + H20 2 Ag + 2 H + ^ 02

the samples were decomposed with a concentrated aqueous KI

solution in a closed pre-evacuated container. The KI solution was acidified with dilute sulfuric acid and saturated with nitrogen gas. The amount of oxygen released was determined

by the weight difference while the I2 was analyzed by standard 117

iodometric titration using thiosulfate

There are obvious flaws in such an analytical method, such as the inaccurate determinations of oxygen released. But such an analysis was simple and allowed a reasonably accurate

estimate of the oxidation state of silver in these compounds. III. FLUOROSULFATES OF SILVER(II)

A. INTRODUCTION

Silver exists in most of its compounds in the +1 oxidation state, corresponding to the electronic configura• tions [Kr] 4d^^. The higher oxidation states, +2 and +3 are far less common. At first sight, the abundance of the univalent silver compounds may be explained by the stability of the closed-shell d1(^ electronic configuration. However, on further examination of other elements in Group IB, the extra stability of the univalent state is unique for silver, with copper(II) and gold(III) most common for these elements.

Hence the stability of a particular oxidation state must involve more than a single contributing factor.

One of the contributing factors would be ionization potentials; the first three for silver and copper are listed in Table 8 for comparison. Even though there are no dramatic differences between them, the balance of all the contributing factors could be offset by the small bias of individual contri• buting factors. The values shown in Table 8 suggest that the

+1 oxidation state of silver is easier to obtain, whereas the divalent state of copper is more favoured, and the +3 oxidation state is easier to obtain for silver based on the sum of the first three ionization potentials.

TABLE 8

SOME IONIZATION POTENTIALS OF SILVER AND COPPER (kJ/mol)

Element First Second Third

Ag [Kr] 4d105s1 730.8 2072 3361

I Ip.Ag = 2803 £ Ip.Ag = 6164 1/2 1/2/3

Cu [Ar] 3d104s1 745.3 1958 3554

I Ip.Cu = 2703 I Ip.Cu = 6257 1/2 1/2/3

From another view point, the reduction potential for the

11 9

Ag(II)/Ag(I) couple is +1.93 V in 4 mol/1 HN03 , whereas the same couple for copper in acidic aqueous solution is only 120

+0.153 V . Hence in aqueous medium, silver(II) is much more difficult to obtain than copper(II) and is a much stronger 66

oxidizing agent. In fact, copper(I) is very unstable in such a medium and readily disproportionates into Cu(II) and metallic copper, where E° for such disproportionation is 120 + 0. 37 V . The stabilities of the silver(I) compounds are further 121 reflected in the very high lattice energies . For example, the value for AgCl is 904 kJ/mol as compared to 775 kJ/mol for NaCl and 703 kJ/mol for KC1, which have similar ionic radii and the same rock salt structure. Hence many silver(I) compounds exist in ionic lattices.

Nevertheless, the higher oxidation states of silver could be made easier to obtain by selecting more favourable reaction medium, such as in bipyridyl complexes, the reduction potential of the Ag(II)/Ag(I) couple is significantly reduced to 1.45 V.

The earlier work on the silver(II) and (III) systems up to 122 123 1960 has been reviewed in detail ' ; and the more recent ones described in the Inorganic Chemistry textbook by Cotton 124 and Wilkinson . During the course of this study, a compre• hensive review on the heterocyclic and macrocyclic complexes 125 of silver(II) and (III) by H.N. Po has appeared

In contrast to the silver(I) system, a lack of simple binary compounds is apparent in the II and III oxidation states.

Although the reaction between elemental silver and fluorine was first described by Moissan in 1891, the existance of silver

difluoride, AgF2, was not confirmed until the 1930's o,i^/^ 67

and has since then remained to be the only higher valent binary compound of silver. A9F2 can ^e obtained as a black microcrystalline solid by the direct fluorination of metallic silver or a number of silver(I) salts such as AgCl. It has been found useful as a versatile oxidizing or fluorinating agent, more reactive than Cobalt(III) fluoride, as in the examples:

12 8

CO + 2 AgF2 • COF2 + 2 AgF

129

S0 F + 2 AgF S02 + 2 AgF2 +~ 2 2

It is also used as a catalyst in catalytic fluorinations ^3^

such as the fluorination of SOF2, to give pentafluorosulfur

131 hypofluorite, SF5OF .

AgF2

SOF2 + 2 F2 • SF5OF 200 °C

132 133

The recently reported crystal structure of AgF2 ' shows an orthorhombic structure with a highly tetragonally distorted octahedral environment for Ag2+. Several magnetic studies

have shown that AgF2 is magnetically concentrated and shows ferromagnetism below 163 K "*"34 136^ The corresponding argentic oxide, AgO, obtained via the oxidation of an alkaline silver(I) solution by peroxydisulfate, _ . 2- 137

S20g , according to:

+ 2 2 2 Ag + S2Og + 4 OH 2 AgO + 2 SO^ + 2 H20

I III has been identified as a mixed valence Ag Ag 02 based on 138 139 its diamagnetism and neutron diffraction studies . The

Ag1 and Ag111 ions in AgO can be chemically separated in an alkaline solution in the presence of complexing agents such

as periodates and tellurates ^•^^/ resulting in the formation of Ag111 complexes as outlined below:

111 4 AgO + 6 KOH + 4 KI04 9*- 2 [Ag (I0g) 2] + Ag20 + H

Other Ag(II) or Ag(III) complexes are summarized and dis• cussed below:

1. Oxysalts of the type Ag(II)(Ag304)2S04 or Ag(Ag304)2X,

where X = N03~, C104~, F~, HS04~. These are obtained by the 122 anodic oxidations of silver(I) salts in aqueous solution

The spontaneous decomposition of these compounds to give AgO according to:

Ag(Ag304)2X AgX + 6 AgO + 0 is accelerated in boiling water "L'i"L.

2. Coordination complexes of heterocyclic and macrocyclic amines of both silver(II) and (III), which have been exten- 125 sively studied and very recently reviewed in detail by Po

These can be separated into: (i) Nitrogen donor ligands such as pyridine, polypyridines, pyrazine of silver(II); macrocycles such as porphyrins and tetraaza complexes of silver(I and (III). (ii) Mixed nitrogen-oxygen donor ligands such as pyridine mono-, di-, and tricarboxylates and pyrazine carboxylates.

In general, several synthetic methods have been applied to prepare these coordination complexes.

(a) By far the most common method is the peroxydisulfate oxidation of a cold aqueous silver(I) solution in the presence of excess ligand molecules.

(b) Anodic oxidation of a silver nitrate solution containing excess ligand in a divided cell is also used.

(c) In rather rare instances where the corresponding silver(I) complex is known and stable to direct oxidation, an oxidizing agent such as ozone or AgO is used.

(d) Direct mixing of silver(I) and certain porphyrin bases and tetraaza ligands may result in a disproportionation to the

142 143 corresponding silver(II) complex and metallic silver ' # All of the above silver(II) complexes are paramagnetic and magnetically dilute with only two proposed cases of 144 antiferromagnetic behaviour in silver(II) bisnicotinate 145 and silver(II) bis(pyrazine) peroxydisulfate Electron spin resonance and electronic spectroscopy, alonT g wit• jh- u X-rav y crystax.l i structuri. e determination, , s 146,147,148 have indicated a square planar coordination around silver in these complexes except the unusual 2,6-pyridinedicarboxylate complexes, which have a highly distorted octahedral struc- 149 150 ture ' . Diamagnetism and hence a spin-paired configura tion has been observed for the silver(III) coordination complexes.

3. Ternary Fluorides of the type:

II IV IV 151 (a) Ag M F6 where M = Ge, Sn, Pb, Ti, Zr, Hf, Pd, Pt

These are prepared by direct fluorination of a stoichiometric mixture of a silver(I) salt and a metal(IV) salt at high temperatures, as in the case of the light blue AgSnF,. shown b below:

F2 Ag„SO. + 2 (NH„)„SnCl, • AgSnF^ 24 4 2 6 45QoC/ 50_7Q hrs> 6

Curie-Weiss behaviour with small Weiss constants have been observed for most except the dark blue-violet hafnium and zirconium complexes, where strong magnetic exchange between 2 + the Ag ions has been suggested. The electronic diffuse reflectance spectra have been fitted to a tetragonally dis- 9 153 torted (elongated) d system using curve fitting techniques

I II III 1 111 (b) M Ag M F6 where M = K, Rb, Cs; M = Al, Ga, In, 152

Tl, Sc, Fe, Co . These are synthesized by direct fluorination of a mixture of alkali metal chloride, silver(I) sulfate and a M(III) salt, usually the oxide, such as A^O^, ^20^,

Ga203, in a 1:1:1 mole ratio at % 550-700 °C. The Tl, In and

Sc complexes have the cubic RbNiCrF, structure and complicated, b not well understood magnetic behaviour. None of the complexes show Curie-Weiss behaviour while all the complexes containing diamagnetic M(III) ions have low magnetic moments between 1.30 and 1.16 y at room temperature.

4. Silver(II) Anionic Fluoro-complexes of the types:

I 1 154 (a) M2 (AgF4) where M = Cs, Rb or K

1 1 155 (b) M (AgF3) where M = Cs, Rb or K

i:C 11 156,157 (c) M (AgF4) where M - Ba, Sr, Ca or Hg

II 1 157 . (d) M2 (AgF6) where M ^ Ba

The first two types are synthesized by heating stoichio• metric mixtures of an alkali fluoride and silver(II) fluoride under an inert atmosphere as shown by the example: 330°C 2 2 CsF + AgF Cs2AgF4 Ar - 12 hrs

The tetrafluoroargentate(II) complexes obey Curie-Weiss law

with magnetic moments of ^ 1.7 yn at room temperature whereas the trifluorocomplexes are magnetically concentrated with muc lower magnetic moments 0.8 y_). B

Complexes of the type (c) and (d) can be prepared by direct fluorination of a mixture of two salts at high tempera ture, in a manner similar to that used to prepare AgMF,. b complexes. For example,

2 BaC03 + Ag2S04 - 2 BaAgF4 500°C - 30 hrs

The high temperature is required to prevent formation of any

silver(III) species. BaAgF4 is found to be isomorphous with

KBrF4 with square planar AgF4 groups; all silver sites are equivalent. Bulk magnetic measurements, electron spin 153 resonance and diffuse reflectance spectra all suggest a 9 tetragonally distorted d environment, in agreement with a silver(II) formulation.

I 5. Anionic Fluoro-complexes of silver(III) as M AgF4 where

1 158 1 I3:i M = Cs or K , and M,M Ag Fc , where M = Cs and 73

M1 = K •LJ:7. The diamagnetic M^AgF^ complexes are obtained

from direct fluorination of a mixture alkali metal chloride or nitrate and silver(I) nitrate between 200 and 400 °c.

However, fluorination of a mixture of CsCl, KCl and silver nitrate in a mole ratio of 2:1:1 at 300 °C for three hours results in a purplish red paramagnetic complex with the com• position Cs2KAgFg. The rather low magnetic moment of 2.6 B.M. is suggested to be due to the presence of decomposition impurities such as CsAgF^, CsF and KF. Analogous copper(III) complexes such as K^CuF^ have been found to have magnetic moments (2.8 B.M.) very close to the spin-only value of 16 0 2.83 y . The presence of paramagnetic silver(III) ions in Cs2KAgFg is supported by the electronic diffuse reflectance 8 161 spectrum which may be assigned to a spin-free d system

In summary, there are no binary silver(II) or (III) compounds except silver difluoride; there are more divalent complexes than trivalent complexes. All silver(II) complexes

9 are paramagnetic, consistent with a 4d electronic configura• tion, while all silver(III) complexes are spin-paired diamag• netic, except Cs2KAgFg. The existence of a large number of donor- and fluoro-complexes are perhaps due to the convenient synthetic routes of either peroxydisulfate oxidation in aqueous medium or direct fluorination of salt mixtures at elevated 74

temperatures.

Silver was chosen as the first transition metal to be investigated in this study for a number of reasons: 16 2

1. It was observed earlier by H.A. Carter in our

laboratory, that the reaction of silver perchlorate, AgC104, with fluorosulfuryl peroxide, S20gF2, produced besides chloryl fluorosulfate, ClC^SO^F, and large amounts of 0^, a black-brown residue with considerable oxidizing ability. A more detailed investigation was postponed as the main interest at the time was the study of chlorine-oxygen compounds.

2. The use of bisfluorosulfuryl peroxide, 820^2, as an oxidizing agent on silver(I) appeared reasonable as the isoelectronic 2- peroxydisulfate ion, S20g , is widely used as the oxidizer in aqueous medium as mentioned earlier.

3. The pseudo-halide character of the fluorosulfates and the abundance of fluoro-complexes of silver(II) and(III) suggest the possibility of corresponding fluorosulfate complexes.

4. The use of silver(II) difluoride as a catalyst in the fluorination of sulfur trioxide to give bisfluorosulfuryl peroxide, 820^2 "*"^3, and fluorine fluorosulfate, FOSC^F , has led to speculations of a silver(II) fluorosulfate as a possible intermediate. 8 5

5. The recent appearance of a paper by Dev and Cady , which describes the reaction of S2°gF2 witn A(?2° an<^ A

silver(II) with the formula Ag20(S03F)2.

Having focused the interest of this study onto silver, a number of possible synthetic routes appear feasible:

(a) The simple oxidative addition of SnO_F to silver(I) z b 2

fluorosulfate to give binary compounds of higher oxidation

states;

(b) The direct insertion of SO^ into AgF-> to give again binary fluorosulfates;

(c) The oxidation of silver metal and a number of silver(I) substrates with suitable oxidizing agents such as S_0_F„ and z b z

BrS03F;

(d) Reacting silver(III) containing AgO with some fluorosulfonating agent to hopefully retain the trivalent state in the product; and

(e) Formation of fluorosulfato complexes analogous to the fluoro-system.

These routes are subsequently attempted and described in the following sections. B. SILVER(II) FLUOROSULFATE

1. introduction

The initial aim was to attempt the straightforward oxidative addition of 820^2 to the commercially available

I Ag S03F to give , hopefully, the corresponding binary silver(II) or even the less likely silver(III) compounds. However, a few unusual and unexpected observations from the first few

S F had creat I reactions with 2°6 2 ed suspicion about this Ag (S03F) reagent. The non-volatile dark brown-black products from reaction at room temperature and at 7 0 °C appeared always non- homogeneous, while the liquid phase from the higher temperature reaction was contaminated with a yellowish volatile by-product.

This was judged to be chlorine fluorosulfate, ClOS02F, assuming

AgCl was present as an impurity in the starting material. The formation of CI2 on interaction with S2°gF2 -*-s likely and the subsequent reaction to give CIOSO2F is well documented -^6^.

The presence of CI2 and CIOSO2F was confirmed by the reaction

of AgCl and Sn0,Fo described later. In addition, the solid Z b Z product formed gave weight increases which indicated a compo•

sition close to Ag(S03F)3. At this time, the infrared spectrum of the starting material in sodium-dried nujol mull showed a complete absence of the normal S0-.F vibrational bands, but two 77

strong bands at 1050 and 600 cm appeared to correspond with

the and of a tetrahedral sulfate group. Finally, this

AgSO^F reagent was found to contain no fluorine at all based

on its elemental analysis performed by Analytische Laboratorien,

West Germany. It was then apparent that the material sold

by Ventron Alfa Corporation as AgSO^F was in reality an impure

Ag2S04 with AgCl as one of its impurities but without even a

trace of fluorine.

The above episode had prompted us to synthesize the real

AgSO.jF (Section II.C. 2) and to extend the study to other

silver(I) salts as well.

2. Synthetic Reactions

(a) Reactions with Bisfluorosulfuryl Peroxide

Bisfluorosulfuryl peroxide was found suitable in oxidizing

a wide variety of silver(I) salts. The reactions performed

are described below,

i) Silver(I) Oxide, Ag20

A sample of Ag,,0 (0.606 g, 5.23 mmol Ag) was introduced

into a preweighed one-part thick-wall pyrex reactor equipped with a teflon coated magnetic stirring bar as described in

Section II.A.3. To eliminate any moisture, the reactor was

flame dried in vacuum and the solid reaqent was handled inside

the drybox. A large excess of S90fiF9 (approximately 10 g) was vacuum distilled into the liquid nitrogen cooled reactor,

the excess So0,F- also served as a working medium. The 2 b 2 mixture was then allowed to warm to room temperature and stirred overnight. Gas evolution was observed. This was assumed to be oxygen, and subsequently pumped off in vacuum.

(Note 1). The dark brown-black solid suspension was then stirred in a water bath (Note 2) at 70 °C for three days. The volatile materials were then distilled off (Note 3) and the remaining dark-brown solid pumped to a constant weight of

1.604 g. To ensure complete reaction, fresh S20gF2 (approxi• mately 5 g) was distilled into the reactor and the mixture again stirred at 70 °C overnight. No further increase in weight was detected after the volatiles were removed. The

(Note 1) The process was to avoid the build up of high pressure inside the reactor. This was particularly necessary for the reaction of silver carbonate, where large amounts of CC^ and gases evolved.

(Note 2) A water bath was used instead of the more convenient oil bath in order to avoid a possible violent explosion or fire if the one-part reactor cracked.

(Note 3) "Bumping" of solid often occurred when distilling volatiles off a suspension. This presented experimental difficulties where the progress of a reaction was followed by weighing the reactor. In addition, solid seated on, the seal of the Teflon valve also allowed slow leakage. Hence removal of volatiles was performed slowly through very small openings to minimize bumping. expected yield for Ag(SO-jF)2 was 1.600 g (5.23 mmol). This hygroscopic product was then analyzed and characterized as described in Section 3 of this chapter. The result of the

elemental analyses is as follow: Calculated for Ag(S03F)2:

%Ag, 35.25; %S, 20.96; %F, 12.42. Found: %Ag, 35.10;

%S, 20.69; %F, 12.15.

This reaction was also carried out at room temperature.

The reaction proceeded much slower as evidenced by the slower

solid product weight increases. The dark brown-black product always appeared to non-homogeneous until finally the weight

ratio of reagent to product corresponded to that of Ag(SC>3F)2

(ii) Reactions of Ag2C03, Ag and AgSG^F

The reactions of silver carbonate, silver metal and silver(I) fluorosulfate were carried out in much the same manner as those of silver(I) oxide. In particular, a more vigorous gas evolution was observed in the reaction of silver carbonate, due to the release of gaseous carbon dioxide. The

reactions of silver metal (Note 4) with S20gF2 proceeded very slowly, and required up to seven days even at 70 °C before completion. In contrast, the laboratory prepared silver(I)

(Note 4) Reactions with silver metal were attempted mainly to avoid any more unpleasant susprises with respect to identity and purity of chemicals. The purity is also the highest among the silver reagents available. 80

fluorosulfate (Section II.C.2) reacted readily at room II

temperature with S206F2 to give Ag (S03F)2 within two days.

All of the above reactions were monitored by frequent

removal of volatiles and determining the up-take in weight until constant weights of non-volatile products were obtained.

The products were identified as Ag(SO.jF)2 by their IR spectra,

the thermal decomposition behaviours (to be discussed in the

next section) and their physical appearances.

(iii) Silver Chloride

The reaction of AgCl with S206F2 was to confirm the

presence of AgCl in the Alfa "AgSO^F" reagent as suggested

from the observations described earlier. Excess S20gF2 (^ 10 g) was vacuum distilled into a one-part thick-wall pyrex reactor

(Section II.A.3.) containing 0.8256 g (5.76 mmol) of AgCl at liquid nitrogen temperature. The mixture was allowed to warm to room temperature and stirred magnetically, the white

AgCl turned gradually brownish while a greenish yellow gas was evolved. Further stirring at 70 °C for three days resulted in a homogeneous dark brown powder suspended in a dark yellow liquid phase. The gas phase infrared spectrum of the 75 volatile material indicated the presence of both S_0,F„ and Z b Z 166

C10S02F . A very small amount of red liquid of very low volatility was also present in the reactor. This was judged

to be chloryl fluorosulfate, C109S0_.F, formed from the hydrolysis of ClOSC^F with residual moisture according to 165

Gilbreath and Cady . Removal of all volatile materials at 50 °C resulted in 1.765 g of a dark brown powder, compared

to 1.763 g (5.76 mmol) expected for Ag(S03F)2. The identity

of the solid product was further confirmed as Ag(S03F)2 by its thermal decomposition and IR spectrum. (iv) Silver Sulfate

The reaction was carried out in direct comparison to

that of the Alfa "AgS03F" reagent. Silver sulfate (0.596 g,

3.82 mmol Ag) and approximately 20 g of S20gF2 were reacted at 70 °C for six days until the weight increases of the solid product ceased, giving 1.100 g of dark brown solid versus

1.169 g expected for Ag(SO^F)2• Small amounts of non-condens• able gas was detected, most likely oxygen released from the sulfate (Note 5). The dark brown powdery product was again

identified to be Ag(S03F)2 by its IR spectrum and thermal

decomposition. The formation of Ag(S03F)2 probably occurs according to the following equation:

70 °C

Ag2S04 + 2 S206F2 »- 2 Ag(S03F)2 + S03 + | 02

(Note 5) The presence of 02 could not be attributed entirely to the reaction of sulfate ions as a noticable portion of S«0,F was also found to decompose to S-CvF- and oxygen at elevated temperatures ( Ref. 125J. (v) Silver(I) Trifluoromethylsulfate and Trifluoroacetate

Reactions of Ag03SCF3 or Ag02CCF3 with S20gF2 were

expected to lead to pure Ag(S03F)2 as the cleavage of the

S-C bond of (03SCF3)~ and C-C bond of (02CCF3)~ should lead

to volatile by-products such as S03, C02 and F3COS02F. Excess

S F waS reacted 2°6 2 with 0.635 g (2.88 mmol) of Ag02CCF3 at

^ 60 °C for three days. The gas phase IR of the volatile

materials indicated the presence of C02 and CF3OS02F

along with So0,Fo. The yield of dark brown solid was found 2 b 2

to be 0.887 g as compared to 0.880 g for Ag(S03F)2. The

product was further identified to be Ag(S03F)2 as in other reactions.

The reaction of AgS03CF3 with excess S20gF2 was carried out at room temperature for six days. During this time, 168 16 9 CF3OS02F and S03 were detected in the gas phase IR spectra of the volatiles. The solid product was identified to

be Ag(S03F)2 by the same means as above. The elemental analysi: of the solid product indicated a slightly impure product.

Calculated for Ag(S03F)2: %Ag, 35.25; %S, 20.96; %F, 12.42.

Found: %Ag, 34.07; %S, 20.50; %F, 12.11.

(vi) Silver Nitrate

Gas evolution was observed when a sample of silver nitrate

(0.972 g, 5.72 mmol) was reacted with approximately 20 g of

S2°6F2 at room.temperature. After removing the non-condensable gas evolved, the reaction mixture was stirred at 60 UC for

six days. The gas phase infrared spectrum showed bands due to S^O^F^ only. The IR spectrum of the non-volatile product 73

indicated the presence of both Ag(SC>3F)2 and NO^SO^F)

The weight of the solid product was approximately 2.62 g, compared to 2.58 g expected for a one-to-one mixture of

Ag(S03F)2 and N02(S03F). (vii) Silver Tetrafluoroborate

AgBF4 (0.968 g, 4.97 mmol) and excess S2OgF2 reacted quite readily with gas evolution at room temperature. The 170 171

presence of BF3 , SiF^ and S20gF2 were detected by the gas phase IR spectrum of the volatiles. After reacting for

24 hours, the weight of non-homogeneous solid was 1.35 g.

Further reaction at 60 °C for five more days increased the yield to a constant value of 1.47 g versus 1.52 g expected for

Ag(S03F)2 and 1.12 g for AgF(S03F). The IR spectrum of the

solid between BaF2 windows showed beside the expected bands

-1 -1 of Ag(S03F)2, shoulders at 1280 cm and 1045 cm , it there• fore appeared that the product was rather impure.

(viii) Silver(I) Fluoride

The incomplete reaction of AgBF^ suggested the possible existence of a mixed fluoride fluorosulfate, hence silver monofluoride (0.823 g, 6.48 mmol) was reacted with excess S2°6F2* The reacti°n appeared to proceed very sluggishly.

After two days at 70 °C, the weight of the black solid particle was 1.017 g, compared to 1.464g expected for AgFfSO^F).

Further reaction appeared to be hindered by initial product coating the surface of the reagent particles. The IR spectrum

of the solid between BaF2 windows showed some weak bands in the S—0 stretching region, perhaps indicative of some fluoro• sulfate group.

(ix) Silver(II) Fluoride

Attempting again to synthesize a mixed fluoride-fluoro-

sulfate, 0.118 g (0.806 mmol) of AgF2 was reacted with excess

S„0rFo at 70 °C, the weight increase of the solid material 2 b 2

after one day was very small, 8 mg against 64 mg for AgF(S03F) and 128 mg for Ag^O-^F),,. Further reaction was probably again hindered by surface coating of starting material.

(b) Reactions with Mixtures of HS0-.F and So0,F„ j 2 6 2 In search of a less time-consuming synthetic route to

Ag(S03F)2 of high purity, the use of HS03F was tested. Brazier 6 6 and Woolf have found silver metal to be soluble in boiling

HSO^F. Recently, a one-to-one by volume mixture of HSO^F and

S20gF2 has also been found useful in synthesizing gold tris- 172 fluorosulfate, Au(S0_jF)3, in our laboratory .• Approximately 3 ml of HSO^F was distilled directly from the acid distil• lation apparatus (Section II.A.4*) into an evacuated one-part pyrex reactor containing 0.293 g (2.72 mmol) of silver powder.

After pumping briefly on the mixture on the vacuum line, about the same volume of ^2^SF2 waS distilled into the reactor at liquid nitrogen temperature. The reactor and contents were allow to warm up to room temperature by themselves. A vigor• ous reaction occured, producing a dark brown precipitate.

After the reaction mixture was stirred for thirty minutes, the unreacted S~0,F„ was removed by vacuum distillation. With the Z b 2 suspension warmed at about 50 °C, the less volatile HSO^F was distilled directly into the safety trap of the vacuum line.

The weight of the dark-brown powery solid was,0.834 g as compa ed to a yield of 0.831 g expected for AgfSO^F^. Further

identification was also made by IR spectroscopy and thermal decomposition.

In some preparations, the HSO^F was more efficiently

removed by vacuum filtration, because of the extremely low

solubility of Ag(S03F)2 in HS03F.

Inside the dry-box, 0.234 g (1.61 mmol) of powdered AgF2 was loaded into a thoroughly dried, preweighed one-part

pyrex reactor. Then a huge excess of SO^ (^ 15 g) was vacuum 86

distilled into the reactor. The mixture was warmed to 50 °C and stirred for thirty minutes. Removal of all excess SO^ at approximately 40 °C resulted in 0.482 g of dark brown solid.

The expected yield from up-take of SO^ to AgtSO^F^ would be

0.492 g. This product was again characterized by its IR spectrum and thermal decomposition.

(d) Reactions with Bromine Monofluorosulfate

The synthetic reactions to Ag(SO_.F)„ via neat S„G>F„ 5 Z Z b Z oxidation were rather slow and in general could be termed as surface reactions due to the poor solvolizing ability of S^O^F^.

This is illustrated by the incomplete reactions of the silver fluorides. Bromine monofluorosulfate has been found to be a better solvating agent and hence may react faster.

(i) Silver Metal, Silver(I) Oxide and Silver(I) Fluorosulfate

An excess of orange-red BrSO^F 5 g) was reacted with silver powder (0.795 g, 7.37 mmol) in a one part thick-wall pyrex reactor at 70 °C for three days. Evolution of reddish gaseous material and the dark red coloration of the solution suggested the formation of bromine. No suspended particles was observed in the dark red BrSO^F solution at 70 °C. Particles appeared to be crystallizing out from solution when cooled to room temperature. Excess BrSO^F and other volatiles were pumped off and 1.76 g of black solid with traces of dark brown 87

material remained. The product weight ratio was less than the expected 2.21 g for AgCSO^F)^. Subsequent stirring in fresh

BrSO^F at 90 °C for one day did not increase the weight of the non-volatile product.

A similar reaction was carried out using A.^*^ as ^e starting material. Except for the evolution of a non-conden• sable gas, presumably oxygen, the observations and results were identical. It was also found that the weights of product yielded per mole of silver used were approximately the same in both reactions and identical IR spectra were obtained. The reactions were repeated at room temperature and at 150 °C while the laboratory prepared Ag^SO^F) was found to react at room temperature. Apart from the speedier dissolution of the reagent at higher reaction temperatures, about 240 g of black solid product per mole of silver in the starting materials were obtained in all cases, independent of the temperature (between

0 and 100 °C) at which the volatiles were removed. The yield of 24 0 g/mole of silver could correspond to a product of the formula Ag^CSO^F)^. Further characterization of this black product will be discussed in Section D of this chapter.

In an attempt to react this black material further, excess

^2°6F2 waS distilled into a reaction vessel containing the non-volatile product of silver metal and BrSO^F. Stirring the mixture at 70 °C for three days resulted in a dark brown powdery solid with a yield corresponding to the value for 88

AgCSO^F)^. The IR spectrum and thermal decomposition further identified the product as AgCSO^F)^. It appears that the black solid from the other reactions could also be converted to AgCSO-jF^ in the same manner.

(ii) Silver(I) Fluoride

As reaction of S„OrF„ with AgF was incomplete, the better 2 o 2 solvolizing, BrSO^F, was attempted. Similar to the other reactions with BrSO^F, AgF was found to react and dissolve in

BrSO^F at room temperature. The gas phase IR spectrum indicated the presence of SiF^, most likely due to the attack of BrF on the pyrex reactor. Removal of all volatiles again yielded

the black solid product. Further reaction with So0,F„ led to 2 o 2 the formation of AgtSO-^F^, identified in the same manner as in the experiments discussed above. The elemental analysis gave values of: %Ag, 35.01; %F, 12.56. Calculated for

Ag(S03F)2: %Ag, 35.25; %F, 12.42.

(e) Discussion

Over the years, a number of attempts to synthesize silver(II) fluorosulfate have appeared in the literature.

However, it appears that the products obtained were either impure or not isolated to be further identified. These obser• vations are summarized as follow:

17 3 (i) As early as 1955, Woolf reported the electrolysis of 89

AgF in HSO^F at approximately 20 mA current. Metallic silver was deposited on the cathode while a black solid was deposited on the anode. The black anode deposit containing 4 0.4 % silver, was assumed to contain divalent silver. The calculated percent silver should be 74.0 in AqF^, 35.2 in AgCSO^F)^ and

47.7 for AgFCSO^F). It seems that the system has not been pursued further.

174

(ii) In 1971, Woolf again reported on the reaction of silver with the BrF^—SO^ system. He stated that the reaction of silver and SO^ in BrF^ was not described in the original I 8 0 synthesis of Ag SO^F as the product was always off-white.

A re-examination of that reaction showed that the product was mainly argentic fluorosulfate which could be thermally decomposed to white argentous fluorosulfate. No further infor• mation was given on the details of the reaction and the product.

However the black anodic deposit from the earlier work on the 173 anodic oxidation of AgF in HSO^F was described to correspond to a 70:30 mixture of Ag(II) and Ag(I).

(iii) . It should also be mentioned that independent observations have been made on the attacks on silver(I) salts by oxidizing agents resulting in the formation of brown solids such as the 162 interactions between S20gF2 and A9C1°4 » BrSO^F and silver 16 2 halide IR windows , and the attack of AgCl by some xenon 175 fluoride-fluorosulfates 90

(iv) An interesting result was found in a recent literature search. Adhami and Herlem conducted a cyclic voltametric studies of iodine, bromine, and sulfur in fluorosulfuric

17 6 acid . Hidden in this paper, is a brief description of the electrochemical behavior of silver. It has been found that the current-voltage plot of Ag+ in HSO^F solution does not give a reduction wave but gives an oxidation wave, which + 2 + represents the oxidation of Ag to Ag . The author postulat• ed that the subsequent reaction was:

2+ + Ag + 2 HS03F —»» AgF2(s) + 2 HSC>3

based on the possible equilibrium

+ HS03F ^ -» HS03 + F"

The oxidation wave occurs at about + 1.3 V, below the oxidation

of the solvent, HS03F, at about + 2 V. A platinum electrode

with an Au(s)/Au(CN)3 reference was used. The insolubility

of AgF2 in HS03F was used to explain the absence of the reduc• tion potential value. But no attempt was made to further identify the precipitate. With hindsight, it may be speculated that the precipitate from the above oxidation process could

well be Ag(S03F)2 rather than AgF2, since the postulated equilibrium is rather unlikely 1~~>. One concluding remark that can be made from this analytical study is that the oxidation of silver to its divalent state by 820^2 should be possible in fluorosulfuric acid as silver is oxidized at a lower potential than the SO^F ion.

All reactions leading to pure silver(II) fluorosulfate, along with their reaction times and temperatures, are summa• rized in Table 9. The most convenient synthetic method to obtain AgtSO^F^ was found to be the oxidation of metallic silver by $20^2 in the presence of HSO^F, requiring a very short reaction time at room temperature. The enhancement of the reaction rate is probably due to the partial solubility of

rr T

metallic silver in HS03F , with the highly soluble Ag (S03F) as a possible intermediate in the oxidation process.

The use of a solution of S2®b^2 ln HS03F as a synthetic reagent is not paralleled in the fluoride system. The use of the corresponding F2—HF system has not been widely uti- 177 lized , mainly due to the corrosiveness of the solution, the below room temperature requirement, and the need for such a synthetic medium when high pressure fluorinations can be employed.

The reaction of Ag20 and S20gF2 described earlier is almost identical to those described by Dev and Cady for the

synthesis of a black Ag^O(SO^F)0 in their note on metal oxy- 92

TABLE 9

SYNTHETIC ROUTES TO Ag(SO_.F)

Starting Oxidizer Time Temp. Reaction Products Material or Reagenta (h) (°C)

Ag S F 168 70 Ag(S03F)2 2°6 2

Ag S„0rF„/HS0oF 0.5 25 Ag(S03F)2 2 o 2 .5

+ Ag2o S F 72 70 2°6 2 Ag(S03F)2 °2

+ + C0 Ag2C03 72 70 S2°6F2 Ag(S03F)2 °2 2

Ag S03F S206F2 48 25 Ag(S03F) 2

AgS03CF3 S206F2 144 25 Ag(S03F)2, CF3S03F, S03

AgC02CF3 S206F2 72 60 CF3OS02F, C02

AgCl S F 72 70 Ag(S03F)2, Cl2, ClOS02F 2°6 2

AgF. SO. 0.5 50 Ag(S03F)2

Ag , AgF BrS03F 48 25-70 Ag3(S03F)4, Br2, BrF, 02

Ag20 or then

AgS03F S F 48 70 Ag(S03F) 2 2°6 2

a. always used in large excess b. 1:1 mixture (by volume) 93

8 5 fluorosulfates , differing only in reaction time and temperature, according to:

+ 25 °C

Ag20 (Ag CO ) + S2OgF2 *• Ag20(S03F)2 (+ C02) 3-12 hrs

Attempts were made to repeat this reaction at the prescribed temperature for periods between three and twelve hours. All such reactions resulted in non-homogeneous black and dark-brown solids. Further oxidation of such at higher temperatures always resulted in the dark-brown AgII(SO F)_. 3 1

It may therefore be concluded that the oxidation of Ag20 or

II Ag2C03 by S^^F-? yields ultimately Ag (S03F)2, presumably

via the black intermediate Ag20(SO.jF)2. Due to the continuous oxidation process with no clearly recognizable intermediate,

isolation of pure Ag20(S03F)2 was not possible. One can

further argue that the evolution of oxygen in the reactions of

Dev and Cady may also arise from the silver(I) oxide or

carbonate, apart from the suggested decomposition of So0^F„ to 2 b 2

^2^5F2 an<^ oxy9en- Furthermore, the infrared spectrum of

the reported Ag20(S03F)2 showed characteristic S—F and S—0

fluorosulfate stretches, but the quality was rather poor that

the band positions were not reported. Based on the known

II reactivity of Ag (S03F)2, it is very likely that their product 94

reacted with nujol and the silver chloride windows used.

Our observation, the conversion of Ag,,0 or Ag2C03 into

Ag(S03F)2, has a similar precedent. The conversion of metal oxides or carbonates into metal fluorosulfates for a number of lanthanides have also been reported more recently by Cady 178 and coworkers

In general, bisfluorosulfuryl peroxide, ^2°6F2' ^S caPa^e of oxidizing silver(I) to silver(II), but a suitable synthetic route results only if the anion can be replaced by a fluoro• sulfate group, thus forming a separable volatile by-product, as in the reaction of AgCl:

70 °C s?°fiF? b 2 AgCl + 2 S2OgF2 w~ 2 Ag(S03F)2 + Cl2 w 2 C10S02F

Otherwise, impure products or reaction mixtures were obtained, as in the case of AgNO^:

2 AgN03 + 3 S206F2 +- 2 Ag(S03F)2 + 2 N02(S03F) + 02

The reaction of AgN0o with S„0,F deserves a further comment.

O Z D Z

The conclusion that resulting solid product was a simple

mixture of Ag(S03F)2 and N02(S03F) rather than a coordination complex was based on the independent presence of vibrational

bands attributable to both Ag(S03F)2 and N02(S03F) in the 95

IR spectrum of the product.

The successful insertion of sulfur trioxide into silver(II) fluoride will be further discussed in Section F of this chapter. The incomplete reactions of the silver(I) and (II) fluorides with S>2®(^2 are PernaPs to ^e expected.

Displacement of fluoride ions in an non-solvating medium would be difficult, in contrast to the complete reaction of

AgF with BrSO^F. However, the reactions showed some up-takes of the fluorosulfate group, indicating the possible existance of a mixed fluoride-fluorosulfate.

After having synthesized silver(II) fluorosulfate via the various routes described, it was of interest to attempt a Born-Haber cycle calculation similar to that for KSO^F

(Appendix B) to estimate the enthalpy of formation of AgfSO-jF^-

The detail calculation is shown in Appendix C. Accordingly,

the AH° for Ag(S03F)2 is found to - 1070 kJ/mol. This value may be compared to the enthalpy of formation for silver difluoride, -210 kJ/mol, obtained from a similar calculation also shown in Appendix C. It should be noted the values obtained from these calculations should be considered as qualitative estimates as mentioned earlier in Section Bl of the introductory chapter. Nevertheless, the conclusions that may be drawn from these calculations are (i) the existance of silver(II) fluorosulfate is thermodynamically favorable, and (ii) the conversion from AgF~ to Ag(SO^F)„ should be feasible. The findings of the synthetic experiments described in the preceding section appear to support these conclusions.

3. Characterization of Ag(SO.jF)2

(a) Infrared Spectra

As already mentioned, the high reactivity of AgCSO^F^

prevented the use of commonly employed infrared window plates

and mulling agents, while inert mulling agents such as halo-

carbon oils are not sufficiently transparent for use in IR

spectroscopy. Hence room temperature infrared spectra were

taken on very thin layers of sample powder pressed between

BaF2 plates. In addition, an IR spectrum at liquid nitrogen

temperature was obtained using a low-temperature cell fitted

with Csl window (Section II.B.l.). The sample was sprayed

onto the central Csl window using nitrogen as a carrier gas.

The reactivity of AgtSO^F^ was sufficiently reduced at ^ 80 K

to prevent window attack. This method allowed observation of

the spectrum down to 200 cm ^, whereas the transparent range of

BaF2 ends at 8 00 cm . It should be mentioned that the region

below 250 cm-"'" was very noisy partly due to adsorbed water on

the window and was also approaching the limit of detectability

of the instrument. Unfortunately, due to the dark brown color

and hence absorption of the incident exciting radiation, no 97

Raman spectrum could be obtained for Ag(S03F)2.

Infrared band positions of Ag(SC>3F)2 and some related compounds are listed in Table 10. As described earlier in the introductory chapter, Section D, the possible coordination modes of the fluorosulfate group can be classified as (la) ionic, (lb) ionic perturbed, (2) covalent monodentate, (3) bidentate, (4) tridentate and (5) tetradentate. From the room temperature IR spectrum, the presence of three dis• tinct, well seperated bands in the S—O stretching region above

1000 cm 1 indicates a loss of symmetry for the fluorosul• fate group. This would eliminate all possibilites except the monodentate, bidentate or perturbed ionic modes. In the low temperature spectrum, no absorption was observed between

8 00 cm 1 and 700 cm ^, where the S—F stretching of ionic fluoro-

47 _i sulfates appear . Furthermore the observed bands 1320 cm ,

1185 cm 1 and 1070 cm 1 fall in the range of bidentate fluoro• sulfates. Agreement is particularly good with the corresponding

1 1 1 76 bands of (CH-j)2 Sn (S03F) at 1350 cm" , 1180 cm" and 1072 cm" . The existance of bridging bidentate fluorosulfate groups in 7 7

(CH3)2Sn(S03F)2 is confirmed by an X-ray diffraction study .

Cu(S03F)2 is reported to have a tetragonally elongated

2+ > octahedral environment for Cu ^•®- f some resemblance is

observed between its IR spectrum (Table 10) and that of Ag(S03F)2, 98

TABLE 10

1 a INFRARED SPECTRA (cm ) OF Ag(S03F)2 AND RELATED COMPOUNDS

Ag(S03F)2 Ag(S03F)2 Cu(S03F)2 (CH3)2Sn- Approx,

(S03F)2 decsrip,

.1320 S,b 1325 ms, sh 1306 vs 1350 s

1310 s

1295 m, sh

1185 vs, b 1205 s, sh 1223 vs 1180 s S03 str region 1185 vs

1070 s, b 1068 s 1115 s 1088 m, sh 1072 s, b

820 ms 837 m 861 s 827 m

824 m SF str 817 m

610 mw 630 m 620 m 592 m 604 m 590 m S03F def modes

557 ms 564 m 554 ms 430 m 428 m 417 s S03F rock modes 416, m 275 m 300 m 304 m

268 ms M-0 str

a. Abbreviations; see Appendix A. b. BaF2 windows, c. At 8 0 K; Csl windows. d. Reference 105.

e. Ref. 76; vibrations due to the Sn(CH3)2 moiety are omitted, 99

It should be noted that the sharpness and resolution of the bands at low temperature were much improved over those at room temperature and small splittings, most likely due to

11 solid state effects, are observed. Since Ag (SO^F)2 is virtually insoluble in HSO^F, and no precedent has been observed for chelating bidentate fluorosulfates, it seems fair to assign a bridging bidentate coordination mode to the fluorosulfates

of Ag(S03F)2>

In summary, strong involvement of two of the three oxygens of the fluorosulfate group in coordination to Ag is indicated with the third oxygen only weakly or not at all coordinated.

This mode is consistent with either a square planar or a tetragonally elongated octahedral environmental for the central atom, resulting in a polymeric structure.

(b) Electronic Spectra

Consistent with the previously discussed vibrational 2+ 9 spectra, two environments for Ag (d ) may be considered as most likely: (i) a square planar, and (ii) a tetragonally 9 distorted octahedral environment. Distortion for a d ion is predicted by the Jahn-Teller theorem. While tetragonal dis- 2+ 2+ tortion is most commonly found for Cu and Ag complexes where structures are reported, with both elongation and 179 compression along the z-axis observed, other distortions , 100

for example, trigonal distortions cannot be ruled out, but they appear to be less common. On the other hand, a square 2+ planar environment for Ag, may be viewed as the extreme case of a tetragonally distorted (elongated) octahedral environment.

The relationship between the spectroscopic terms arised from the Russell-Saunders coupling scheme and the electronic configuration of a silver(II) ion is shown in Figure 9. Depen• ding on the extent of the tetragonal distortion, the energy 2 of the d orbital mayJ fall below the rising value of d orbital z xy towards the extreme, a square planar geometry; in such a case 2 2 the relative ordering of the A, and B„ term would then lg 2g be inverted. Three spin-allowed transitions are expected based on Figure 9. but the exact ordering of the transitions is

uncertain, due to the uncertainty in the ordering of the d xy 2 and d orbitals, z

The electronic spectral data of Ag(S03F)2 along with those

of AgF2 and Cu(S03F)2 are listed in Table 11. Due to the

extremely low solubility of Ag(S03F)2 in HSO-jF, a saturated but very dilute solution showed a single very broad band at 3 -1 approximately 25 x 10 cm . Better resolution was obtained in the diffuse reflectance spectrum. In addition to an intense 3 -1 UV band above 24 x 10 cm , most likely due to charge-transfer, 101

FIGURE 9

ELECTRONIC CONFIGURATION AND SPECTROSCOPIC TERMS OF SILVER(II)

OCTAHEDRAL AND ELONGATED TETRAGONAL LIGAND FIELD

d 2 2 (br ) x -y lg' g g

2 1 a < lg> 2g 2g'

(a (b2g> Xg lg>

d , d (e ) b xz yz v g' Sg

Free Octahedral Elongated Free Octahedral Elongated ion tetragonal ion

0, D 4h 0, D 4h 102

TABLE 11

ELECTRONIC SPECTRA OF Ag(S03F)2 AND RELATED COMPOUNDS

Compound Type of Spectrum X (cm-1 x 103) Ref max

Ag(S03F)2 in HS03F %25 this work

Ag(S03F)2 diff. reflect. >25 vs this work 22.0 a,b 16.6 sh 14.1 sh

AgF, diff. reflect, 22.0 s,b this work 18.0 sh 13.7 w

Cu(S03F) Mull 10.4 (105) 103

three other bands are observed. Tentative assignment may be 3-12 2 made for the broad band at 22.0 x 10 cm as B, —* E ; lg g' 3-12 2 the weaker shoulder at 16.6 x 10 cm as B, —> B~ and lg 2g; 3 -1 the broad band at lowest frequency, 14.1 x 10 cm as the 2 2 2 2 B, —> A, transition. The transition between B, —> E is lg lg lg g expected to be split by spin-orbit coupling. However, in view of the rather poor quality of the spectrum, such splitting is not likely to be resolved. To a first approximation, the ligand field parameter, 10 Dq, may be given 16,600 cm ^. 3 -1

A broad band at approximately 10.4 x 10 cm m the mull spectrum of CufSO^F),, has been attributed to various unresolved d—d transitions in an elongated octahedral complex of D^ symmetry Besides the expected band shifts to higher energy due to the increase in ligand field parameter, Dq, with increa• sing principal quantum number, the electronic spectra of silver(II) should be analogous to the corresponding copper(II) in an identical or similar chemical and geometrical environment.

In view of the similarities between fluoride and fluoro• sulfate in regard to ligand field strength, the diffuse

reflectance spectrum of the commercially available AgF2 was also recorded. An orthorhombic structure with a Pbca space group 132 133

has been determined for AgF2 from neutron powder data ' , with the silver ions surrounded by six fluorine atoms in a four short (average 207 pm) and two long (average ,258 pm) manner corresponding to an extremely distorted (elongated) octahedron.

Details of its diffuse reflectance spectrum are listed in

Table 11. The similarity to that of Ag^O^F^ is apparent. 3 -1 While the intense band above 25 x 10 cm is absent, the 3-1 3-1 three weaker bands appear at 22.0 x 10 cm , 18.0 x 10 cm 3-1 as a shoulder, and 13.7 x 10 cm . Hence the result of the

electronic spectra of Ag(S03F)2 is consistent with a highly tetragonally distorted (elongated) octahedral arrangement around Ag .

Paramagnetism arises from the spin and the orbital angualr momentum of ions. Orbital angular momentum may be associated with the ability to rotate one orbital about an axis to give an identical and degenerate orbital. The orbital angular momentum contribution to the observed paramagnetism would require a three fold degenerate ground state. As the result of the preceding discussion, the most likely ground state for 2+ 2

Ag is ^ig'nenc e an orbitally non-degenerate state. Without orbital contribution, the magnetic moment of octahedral silver(II) complexes should be close to the spin-only value of 1.73 u_.

The magnetic moments of magnetically dilute silver(II) complexes obtained experimentally are usually in the range of 105

1.80 - 2.20 y ^ v, J. _> ^ appreciably higher than the spin-only value. The increase above the spin-only value arises from spin-orbit coupling, where some of the orbitally degenerate excited state is "mixed-in". The ground term may therefore aquire some orbital angular momentum. Nevertheless, measure• ments of the magnetic moment of Ag^CSO^F^ should provide 9 evidence of a paramagnetic d ion and the extent of state mixing due to spin-orbit coupling may be estimated using the equation

y = (1 eff " 2X/|10 Dq|) yg (3.1)

where y , = spin-only magnetic moment *s .o. 10 Dq = the ligand field splitting parameter

X spin-orbit coupling constant

Strictly speaking, equation (3.1) applies to a chromophore with 0.^ symmetry only. However, the distortion from 0^ in this case may be small enough that, to a first approximation, equation (3.1) could be used. The magnetic susceptibility of

Ag(S0^F)2 was measured between 300 K and 80 K by the Gouy method.

The detailed calculation method to obtain the effective magnetic moment is shown in Appendix D. The magnetic susceptibility and moment data are listed in Table 12. A plot of the inverse 106

TABLE 12

MAGNETIC SUSCEPTIBILITIES AND MAGNETIC MOMENTS OF Ag(SO-.F)

yeff (K) (106 cm3 mol 1) (yn)

301 1631 1.91

276 1792 1.92

249 2013 1.92

224 2248 1.92

200 2566 1.92

175 2959 1. 91

149 3540 1.91

128 4217 1.91

105 5440 1.92

80 7574 1.90 a. Magnetic moments are calculated by using the Curie-Weiss law:

r 1/2 u ff = 2.828[x^° (T - 0)] 107

of molar magnetic susceptibility versus temperature is shown in Figure 10. In the same figure is a similar plot for the complex [Ag(bipy)2](SO^F)2/ which will be discussed in the next section.

A linear relationship is obtained from the plot for

AqiSO^F)^, with an intercept of + 20.3 K at the temperature axis. Hence the Curie-Weiss law is obeyed and can be express• ed as:

r 2 yeff = 2.828[XJJ° (T - 0)] (3.2)

where 0 is the Weiss constant. The effective magnetic moment,

ye^^ at room temperature is determined to be 1.92 y^. A temperature independent moment of 1.91 ± 0.01 y_, over the range of 301 K and 80 K is calculated from equation (3.2) .

Using the previously estimated 10 Dq value of 16,600 cm 1 ob• tained from the electronic spectra, the spin-orbit coupling constant, X, is calculated to be - 910 cm 1 using equation (3.1).

This represents a substantial reduction of X below the free-ion

—1 2+ 181 value of approximately - 1840 cm for Ag . This reduction is presumably due to the delocalization of electron density

out of the t2 orbitals of the silver(II) ion onto the fluoro• sulfate ligands.

The Curie-Weiss behaviour of Ag(S0^F)o with a small Weiss FIGURE 10

corr

0 100 200 300 Temp (°K) constant indicates a magnetically dilute system which is in

strong contrast to the findings for AgF2, the subjects of

several studies "*"34 136^ silver(II) fluoride is magnetical•

ly concentrated and shows ferromagnetism below 16 3 K . The

effective magnetic moment at room temperature is 2.0 ± 0.1 u B

and shows Curie behaviour at higher temperatures. The measured magnetic moment of AgCSO^F^ was magnetic field

independent, indicating an absence of ferromagnetic interactions

It appears that the bulkier SO^F groups prevent magnetic ex- 9 0 change far more efficiently than monoatomic halide ligands ,

62 105 as in the previously studied Cu(S03F)2 ' > Copper(II) fluorosulfate shows Curie-Weiss behaviour between approximately 310 and 90 K with a small Weiss constant 62 6 2 of - 2 K . The room temperature magnetic moment of 2.05 y B or 2.08 y_ is independent of temperature. The magnetic moment of silver(II) fluorosulfate is smaller than that

of Cu(S03F)2. Assuming a similar environment for the metal ions in both compounds and a reduction of 10 Dq by the customary 2+ 2 + 30 to 5 0 % when going from Ag to Cu , the lower Veff values for Ag(S03F)2 point to a larger reduction of A below the free 2+ 2 + ion value for Ag then for Cu

The result of the magnetic measurements of Ag(S03F)2 and some related compounds are summarized in Table 13 for comparison

Hence Ag(S03F)2 is paramagnetic with a single unpaired electron TABLE 13

' MAGNETIC PROPERTIES OF Ag(S03F) AND RELATED COMPOUNDS

298 Compound Temp. Range y Weiss Constant Comment Ref. eff

K y B 0(K)

Ag(S03F)2 80 - 301 1.92 + 20 Curie-Weiss this work

AgF2 8 - 500 2.0 Ferromagnetic <163 K 135

Cu(S03F)2 100 - 312 2.08 Curie-Weiss 105

Cu(S03F)2 90 - 298 2.05 - 2 Curie-Weiss 62 9 corresponding to a d electronic configuration. The temperature 2 independent magnetic moment is consistent with a B, ground

state of a tetragonally elongated octahedral or square planar

environment, as suggested from the IR and electronic spectra discussed previously.

(d) Electron Spin Resonance Spectra

The electron possess a magnetic moment and may be treated as a bar magnet. In the presence of an external magnetic

field, free electrons will either align with the magnetic

field, resulting in the lower energy state; or in the opposite direction as the magnetic field, resulting in the higher energy state. This effect is known as the first-order Zeeman effect.

The transition between the two states occcurs upon absorption of radiation in the microwave region. The change in energy is given by:

AE = hv = g 6H (3.3) ^e o

where h is the Planck's constant,

v is the frequency of radiation,

H the applied magnetic field, o is

6 is the Bohr Magneton, and

is the Lande splitting factor. The situation with silver(II) ions is in fact similar; 9 the 4d electronic configuration has an effective spin (S)

of 1/2, and spin angular momentum of Mg = ± 1/2, which is identical to a free electron. However, the Lande splitting factor g depends on the orientation of the molecule with respect to the magnetic field. The g tensor will be isotropic, in other words, independent of orientation if the silver(II) ion is located in a perfectly cubic crystal site with for example, 0^ symmetry. If the symmetry of the crystal site is reduced, the g factor is orientation dependent, that is, anisotropic. But when powdered samples are examined, isotropic spectrum may result from compounds with reduced symmetry due to the overall self-cancelling of the grossly misaligned sets of symmetry axes. It is conventional to define the z-direction to align with the rotational axis of highest order and the associated g factor as g or g... While the g components z I I perpendicular to the external magnetic field be g | = g^. = g^.

It is obvious for a regular octahedral system, q = q = g -3 x y z

= gj_ = g( | = gQ. Another significance of the g factor is the deviation from the free electron value, g , which reflects the amount ^e of orbital contribution. As discussed in the magnetic suscepti• bility section, the orbital angular momentum is not expected 2 with the groundstate B^ . The g-value should be the same as 113

that of a free electron, gg, which is 2.0023. However, the effect of spin-orbit coupling would alter the value of the

g-tensor. Hence equation (3.3) should be substituted by equa•

tion (3.4) in a complex ion.

AE = hv = goBHQ (3.4) or

go = (3-5) m o

where gQ is the g-tensor determined experimentally. It is also the root-mean-square average of the g components in an

anisotropic spectrum.

go = [ I ( gz + gy + gx ^ 1 2 <3'6> ?

This gQ tensor is connected to the spin-orbital coupling constant X and the Ligand field splitting parameter 10 Dq according to:

1 Pi n g ±OU = 2.00 ( 1 - 2 A/| 10 Dq | ) (3.7)

Substituting equation (3.7) into equation (3.1) results in a relationship that connects the effective magnetic moment with g according to: ^o ^

r which is equivalent to:

y = s s + 1 2 3 9 eff got ( )] ( - ) where S = effective spin of the system = 1/2 for Ag(II).

In addition, a nucleus with a non-zero nuclear spin I, will interact with the unpaired electron in the magnetic field to split the absorption into 21+1 components. This is known as the nuclear spin-electron spin hyperfine coupling 107 109

Since both natural isotopes of silver, Ag and Ag, have a nuclear spin I = 1/2, such hyperfine coupling would result in a two identical line spectrum.

The E.S.R. spectrum of a powdered sample of Ag(S03F)2 measured at room temperature showed a single isotropic broad line which only sharpened slightly at liquid nitrogen tempera

ture. No hyperfine structure was detected. The isotropic

spectrum is most likely due to misaligned sets of axes rather than to a regular octahedral or tetrahedral symmetry around 2+ the Ag ions. The isotropic spectrum resolved into a three

component anisotropic spectrum when a solution of Ag(S03F)2

in BrSO-jF was measured at 80 K. The data obtained from these

spectra are listed in Table 14. Sample spectra are shown in

Figure 11. 115

FIGURE 11

ESR-SPECTRA ot 80°K a) solid Ag^SCLF). 3' '2

g0 =2.187

b) Ag^SCy^ in BrSO^F

g,= g„ = 2.407 g2 = 2.096

g3 = 2.072

(g0 = 2.193) 116

TABLE 14

ESR DATA OF SILVER(II) FLUOROSULFATE

Physical State T g^ g, (g ) g0 g- V^*^ V, >o ylvyz' y2 ^3 Meff Meff

( K ) ( yB )

solid powder 295 2.220 1.92 1.91 solid powder 80 2.187 1.89 1.90

BrS03F solution 80 2.198 2.407 2.096 2.077 1.90

a. calculated using equation (3.9)

b. determined by the Gouy method (Table 12)

was The effective magnetic moment, Veff» calculated from

the gQ value using equation (3.9), very good agreement is obtained with the values measured by the Gouy method. The good agreement between the values calculated from the ESR spectrum of the BrSO^F solution and from bulk solid measure• ment suggests little change in the complex on dissolution, besides the reduction in line broadening and the improved resolution of an anisotropic spectrum. 117

(e) Oxidation State Determination

The experimental method has been described in the

Experimental Chapter, Section II.D. 2. A sample of AgCSO^F),,, 2 +

0.8841 g (2.889 mmol Ag ) was decomposed in concentrated aqueous potassium iodide solution saturated with nitrogen gas. The amount of uncondensable gas evolved was determined by weight difference to be 0.0092 g or 0.29 mmol of 0^• The amount of liberated was titrated by standarized thiosulfate solution to be 0.835 mmol according to:

2 2 I2 + 2 S203 — 2 1 + S4Og (3.10)

Assuming the production of 02 and 1^ arose from

2+ + + 2 Ag + H20 »- 2 Ag + 2 H + | 02 (3.11) and

Ag2+ + 2 KI +- Agl + i I + 2 K+ (3.12)

2 +

The amount of Ag present in the original sample would be 2.83 mmol as compared to 2.889 mmol theoretically. This gives an overall oxidation state of approximately 1.98 for the sample. 118

(f) Thermal Decomposition

Under a sealed nitrogen atmosphere, AgCSO^F)^ was found to be thermally stable up to + 210 °C, at which temperature, a dark yellowish gas evolved and condensed to a colorless liquid in the colder part of the vessel. The solid material melted to a black liquid and solidified on cooling to a black residue.

To further investigate the decomposition, a sample of

AgCSO^F), 1.420 g(4.64 mmol), contained in an evacuated one part thick wall pyrex reactor was heated in an oil bath at

215 °C for one hour. Connected to the sample container was another one-part trap cooled under liquid nitrogen to trap the volatile materials. The gas evolved condensed to a white solid in the cold trap while no uncondensable material was detected. The gas-phase IR of the volatiles evolved indicated

75 17 0 the presence of S2°gF2 and S:*-F4 (Note 6). Hence it appears that the initial dark yellowish gas evolved was the fluorosulfate radical. The greyish white residue was identified I 8 2 by its Infrared spectrum to be Ag (SO^F) . If the decom• position is formulated as: 250 °C

S F 2 Ag(S03F)2 »- 2 AgS03F + 2°6 2 (3.13)

(Note 6) SiF^ is a commonly encountered impurity when fluorine containing compounds are heated in a glass container. 119

The expected amount of S„0,F_ produced would be 0.459 g as 2 b 2 compared to the 0.444 g (2.24 mmol) trapped, while the weight of the residue should be 0.961 g versus the 0.976 g which remained.

The formation of bisfluorosulfuryl peroxide by pyrolysis 16 of a metal fluorosulfate is unprecedented . Recently, such

F a pyrolysis formation of $2®6 2 ^as a^so been observed for 18 2 another noble metal, palladium , in our laboratory accord• ing to:

II IV 160 °c Pd Pd (SO.F), +~ 2 Pd(SO-,F)_ + So0cF_ (3.14) 3 6 3 2 2 b 2

C. COORDINATION COMPLEX OF Ag(S03F)2, [Ag(bipy)2](S03F)2

1. Introduction

As mentioned in Section A of this chapter, there are numerous known examples where silver(II) ions are stabilized by nitrogen donor ligands such as 2,21-bipyridine and 122 123 125

1,10-phenanthroline ' ' , generally obtained, when silver is oxidized in the presence of suitable ligands. The

conversion of Ag(S03F)2 into a bis(2,2'-bipyridyl) bisfluoro-

sulfate complex would be a chemical identification of the 120

existence of AgiT ions and would represent an alternate route to these compounds. The ligand 2,21-bipyridine was

2+ chosen because the cation [Ag(bipy)2] has been studied extensively by magnetic susceptibility measurements, electro- 122 123 125 nic and ESR spectra ' ' .In addition, two X-ray

diffraction studies have also been published on [Ag(bipy)2]-

147 148 (N03)2 and [Ag(bipy)2](N03)2*H20 . Furthermore, the

corresponding trifluoromethylsulfate complex, [Ag(bipy)2]- 113

(S03CF3)2 is kncvn and should provide interesting compari• sons. Hence this trifluoromethylsulfate complex was also 113 synthesized according to the literature method for direct comparison (Section II.C.2.). 2. Synthesis of Silver(II) Bis(2,2'-bipyridyl) Bis(fluorosulfate)

In a typical preparation, 0.894 g (2.92 mmol) of Ag(S03F)2 was added under a dry nitrogen atmosphere to a concentrated solution of excess 2,2'-bipyridine (1.110 g, 7.1 mmol) in dry acetonitrile. This was accomplished by loading the sample of

Ag(S03F)2 into the inlet arm of the two-part reactor (Section II.

A.I.). After removing the apparatus from the drybox, the acetonitrile solution was first cooled to - 40 °C before mixing the reagents. The resulting brick red mixture was allowed to warm to room temperature and stirred for a further thirty minutes. A brick red solid separated from the brownish red solution and was vacuum filtered. The product was washed under dry nitrogen with 5 ml portions of dry dichloromethane and subsequently dried in vacuo. Approximately

1.34 g of brick red powder was recovered, corresponding to

a 74 % yield based on the amount of Ag(SO.jF)2 used. This product was found to melt between 218 - 219 °C while decompo• sing to a black liquid. The elemental analysis was found to

S0 F AS F WS: correspond to Ag (C^QHQN,,)2 ( 3 ) 2 °H° Calculated,

%Ag, 17.44; %F, 6.14; %C, 38.85; %N, 9.06; %H, 2.61. Found,

%Ag, 17.15; %F, 6.29; %C, 38.73; %N, 9.01; %H, 2.76.

3. Characterizations

(a) Vibrational Spectra

[Ag(bipy)2](SO^F)2 was found to be much less reactive towards IR window materials and hence a nujol mull spectrum between KRS-5 windows was obtained and shown in Figure 12.

However, only a poor quality Raman spectrum could be recorded, perhaps due to rather dark color of the compound. The vibrational band position are listed in Table 15 along with

the IR frequencies of the corresponding [Ag(bipy)2](CF^SO^)2 complex for direct comparison.

Although two independent elaborate correlation studies of infrared spectra of large groups of metal 2,2'-bipyridine FIGURE 12

The Infrored Spectrum of [Ag(dipy?](SO^F ) between 1600 and 400 cm-'

1600 1400 1200 1000 800

N denotes Nujol A denotes Anion bands TABLE 15 123

3 VIBRATIONAL SPECTRA OF [Ag(bipy)2] (S03F)2 AND [Ag(bipy)2] (CF3S03F)

[Ag(bipy) 2] (S03F)2 [Ag(bipy) ] (CF3S03F)2

Raman IR IRC

1600 s 1602 s 1600 m 1575 w 1568 ms 1570 w 1504 w 1498 s 1497 m 1475 s 1475 m 1445 s 1445 s 1320 m, sh 1320 m 1324 m 1302 vs, b * 1280 s, sh * 1270 vs, b * 1255 vs, b * 1250 mw, sh 1255 s * 1178 w 1175 w, sh 1157 m 1145 vs * 1105 m 1107 mw 1074 m 1075 vs * 1060 w 1033 s 1028 m 1030 vs * 1017 mw 970 vw 970 vw, sh 903 w 902 w 818 vw 820 vw 782 s 782 s 770 s 770 s 760 m * 730 s * 725 s, sh 725 s 663 s 660 m 660 m 651 w 650 w 645 m 635 vs * 580 vs * 573 s * 561 m * 520 s * 440 w 440 vw 418 w 415 ms * 417 s * 408 w 363 s 355 vw 360 w 350 w

* anion bands

a. in cm \ see Appendix A for abbreviations b. nujol and hexachlorobutadiene mull between KRS-5 plates c. neat powder between KRS-5 plates 12.4-

18 3 184

and 1,1O-phenanthroline complexes have been made ' , no

detail assignment of the vibration bands of 2,2'-bipyridine complexes have been published. To attempt such an assignment was considered not within the scope of this study. In any

case, the spectra of the fluorosulfate and the trifluoromethylsulfate complexes are virtually identical except for

the rather intense anion bands. This observation agrees with 18 3 184 the findings of the correlation studies ' where in fact, the spectra obtained are similar in gross features and no obvious correlations with magnetic or other physical properties can be made.

The assignment for the anion bands of [Ag(bipy)^1(SO^F)^ is shown in Table 16 and compared to those of KSO^F and AgSO^F.

Non-coordinated fluorosulfate ions with C^y symmetry appears to best represent the coordination mode, even though the degenerate asymmetric SO^ stretching mode is split by approxi• mately thirty wavenumbers, presumably due to site-symmetry effects. The band positions agree well with reports on ionic *i i * *. 45, 64,105,82 , . _ fluorosulfate compounds , such as the examples of

KS03F and AgSC^F shown in Table 16. The IR spectrum of AgSO^F reported here agrees well with 8 2 the. original reference . In some instances better resolution was obtained as compared to the spectrum recorded during the mid 1950's. As in the original reference, the present IR . 125

TABLE 16

ANIONIC INFRARED BANDS OF [Ag(bipy)^\ (S03F)2 AND RELATED COMPOUNDS

62 [Ag(bipy)2](S03F)2 KS03F AgS03F Assignment

1302 vs 1310 s,b,sh 1280 s v4 (E) S03 asym. str. 1270 vs 1250 s,b

1075 vs 1075 s,sh 1080 s v1(A1) S03 sym. str, 1062 s

730 s 795 s 750 s v2(A1) SF str, 748 s

580 vs 595 s,sh 590 s v5(E) S03 asym. def, 587 s

561 m 577 s 570 sh v3(A1) S03 Sym. def, 561. s,sh

415 ms 410 vs 408 m,b v,(E) SO_F rock 6 3 395 vw 126

spectrum shows ionic fluorosulfate groups with C^v symmetry.

However, all six fundamentals are doubled. This is best

explained by assuming two independent sites for the anion

in the solid state, as in the previously reported Sr(S03F)2

(b) Electronic Spectra

The stability of silver(II) in the cation of [Ag(bipy)2]-

(SO^F)2 i-s further demonstrated by the visible-ultraviolet

spectrum of the complex in aqueous solution. An essentially

identical spectrum was obtained for the trifluoromethylsulfate

complex synthesized here, which also agrees well with the original reference 113. Solid state diffuse reflectance

spectra of the two complexes were also recorded. All data from these electronic spectra along with some relevant references are listed in Table 17. In general, the occurrence of a broad band at approximately 22,000 cm 1 is regarded as evidence for a square-planar environment in nitrogen donor coordination 125 complexes of silver(II) . Such a broad band is likely to be the unresolved combination of the d—d transition discussed in

Section B.3(b) of this Chapter. In the diffuse reflectance 18 5

spectrum of the analogous [Cu(py)4] (S03F)2 , a broad band centered at 17,000 cm 1 is attributed to the unresolved d—d transitions. The position of the band maximum is approximately

30 % lower in energy than that of [Ag (dipy)2 ] (S03F)i2 , which is TABLE 17

ELECTRONIC SPECTRA OF [Ag(bipy)2](S03F)2 AND RELATED COMPOUNDS

-1 3 Compound Type of X , cm x 10 Ref, max _^ _^ Spectrum (e , M cm ) max .

[Ag(bipy)2](S03F)2 in H20 22 (1630) This work 35.7 (22000) 42.9 (16300)

[Ag(bipy)2](S03CF3)2 in H2 0 22 (2160) 113, 35.7 this work 42.9

[Ag(bipy)2](N03)2 in H20 22 (2160) 125,187

]S in H20 22 (1600) [Ag(bipy) 2 2°8 186 28 (very intense)

[Cu(py)4](S03F)2 Diff. refl 17.0 185

[Ag(bipy)21(S03F)2 Diff. refl >28, 22 This work

[Ag(bipy)2](CF3S03)2 Diff. refl 22 This work

[Ag (bipy) 2] (NO-j) 2 in H20 22, 35.7, 42.9 113 128:

not unreasonable considering the decrease in ligand field splitting going from the 4d to 3d series. The rather high extinction coefficient (e = 1630 M "'"cm ^) of the broad band at

22,000 cm most likely arises from intensity stealing of the neighboring huge charge-transfer absorption at 35,700 cm ^

(e = 22,000) .

The magnetic behaviour of [Ag(bipy)2](SO^F)2 was studied between 305 and 80 K with the data listed in Table 18. The cor

plot of l/xM versus temperature has already been shown in

Figure 10 along with the plot for AgCSO^F^. Curie-Weiss behaviour is observed in the temperature range studied with a

Weiss constant of + 5.8 K. The effective magnetic moment is temperature independent at a value of 1.8 2 ± 0.01 y . Agree- B 1 2—188 ment with older work on [Ag(bipy)„]X„ complexes (X = ^-So0o , 2 2 2 2 o

188 113 188 188 113 N03" ' f CIO3" , C104" , CF3S03" ) is rather poor except for the peroxydisulfate complex where an effective

magnetic moment of 1.80 yB is reported. No detailed temperature dependence measurements are reported in these studies and

room temperature measurements on [Ag(bipy)2](S03CF3)2 by Thorpe y r:f values are found between 2.08 and 2.29 y_,. In particular,

3 and Kochi gives a xM value of 2551 x 10^ cm mol which

would yield a Meff value of ^ 2.5 yB- However, the ESR spectrum 129

TABLE 18

MAGNETIC SUSCEPTIBILITIES AND MAGNETIC MOMENTS OF [Ag(bipy)2](S03F)

COr a T X y xm Meff (K) (106 cm3 mol"1) (u_) B

305 1393 1.83

280 1507 1.82

255 1665 1.82

232 1850 1.83

203 2122 1.83

187 2317 1.83

154 2796 1.82

129 3319 1.81

109 4087 1.83

80 5606 1.82

Magnetic moments are calculated by using the Curie-Weiss law:

r 1/2 y^ff = 2.828[XS° (T - 0)] 130

recorded on the sample synthesized in this study gives no

indication of an unusually high magnetic moment for this

compound. Further discussion on the ESR spectrum will be

made later on.

On the other hand, the observed temperature independence,

the small Weiss constant, and the y ff found for [Ag(bipy)2]-

(SO^F)2 agree well with the results of magnetic measurements 18 9 on a series of silver(II) salts of pyridine-carboxylic acids with the exception of silver(II) nicotinate, which was found to 189 144

be antiferromagnetic ' . The effective magnetic moment

of these complexes range from 1.78 to 1.8 2 y . ,

(d) ESR Spectra

Powdered samples of [Ag(bipy)^] (SO^F)2 and [Ag(bipy)21-

(CF3S03)2 were studied at both room temperature and at 80 K.

Since both complexes are slightly soluble in acetonitrile,

ESR spectra were also recorded on saturated solutions of each

of them at 80 K. The spectra obtained were all anisotropic

with two distinct components of the g-tensor. Nuclear-spin

hyperfine coupling was not observed in any case. The data

obtained is tabulated in Table 19. Agreement with previous 2+ 190 191 work on complexes containing the [Agtbipy^l ion ' is

good. Both the fluorosulfate and the trifluoromethylsulfate

complexes give virtually identical g values in the solid state TABLE 19

ESR DATA OF SOME 2,2•-BIPYRIDINE COMPLEXES OF SILVER(II)

Compound T, K g V o g P ll 9i eff* B Ref

[Ag(bipy)2](S03F)2 (s) 295 2.099 2.178 2.058 1.82 Thi-s work

(bipy)2](S03F)2 (s) 80 [Ag 2.091 2.170 2.051 1.81 This work

[Ag(bipy)2](S03F)2 in CH3CN 80 2.092 2.166 2.054 1.81 This work

(bipy)2](CF3S03)2 (s) [Ag 80 2.087 2.164 2.047 1.81 This work

[Ag(bipy)2](CF3S03)2 80 2.092 2.160 2.057 1.81 This work

(in CH3CN)

tAg(bipy)2](S2Og)(s) 77 2.094 2.184 2.047 1.81 191

[Ag(bipy)2](C104)2 (s) 77 2.087 2.169 2.045 1.81 190

a. Calculated from equation (3.9) and in frozen solutions of CH^CN. Good agreement is also obtained between the magnetic moments calculated from the g values and those from bulk susceptibility measurements.

For [Ag(bipy)2](CF^SO^)2/ the calculated Ueff is 1.81, iden• tical to those of the corresponding fluorosulfate complex.

In summary, the complex [Ag(bipy)^] (SO^F)^ appears to contain a silver(II) ion coordinated with two 2,2'-bipyridine ligands in a square planar fashion while the flurorsulfate anions are ionic in nature with very weak or no interaction with the silver ion.

D. ANIONIC FLUOROSULFATO COMPLEXES OF SILVER(II)

1. Introduction

As mentioned in Section B of this chapter, a black solid product was obtained in reactions involving BrOS02F as fluoro-

sulfonating agent. The weight ratio of product to reagent

suggested the formula Ag^SO^Fj^. The oxidizing ability of this product suggested the presence of silver in its higher oxidation states. A reasonable formulation would be

Ag^Ag11(SO^F)^ suggesting perhaps the presence of the anion 112- [Ag (SO..F).] . Ag„ [Ag(SO^F) . ] would be the first example of 133

a mixed-valence state Ag(I)—Ag(II) compound. Previously 192 I in reported examples involve the valence pairs Ag —Ag such

I III 1 as Ag [Ag 02] or Ag°—Ag as in the silver subfluoride Ag2F.

The formulation of the mixed valency compounds as

C 11 Ag2" [Ag (SO^F) ^ ] suggests the possibility of synthesizing analogous complexes where Ag1 is replaced by other univalent

cations. The synthesis of the complex K2[Ag(SO^F)^] is des• cribed in the following section. The fluoro-analogue of this 154 15 5 complex has been reported by Hoppe, and coworkers ' in the group of complexes M^AgF^, where M1 = K, Rb, or Cs.

2. Syntheses And Elemental Analyses

(a) Ag3(S03F)4

Reaction details leading to the formation of the hygro• scopic trisilver tetrakisfluorosulfate have been described in

Section B.2(d) of this chapter. Elemental analysis:

Calculated for Ag3(S03F)4: %Ag, 44.95; %S, 17.82; %F, 10.56.

Found: %Ag, 44.85; %S, 17.78; %F, 10.45.

It is interesting to note that the black product described 17 3

by Woolf in the electrolytic oxidation of AgF in HS03F appeared similar but only contained 4 0.4 % Ag.

(b) K2[Ag(S03F)4]

Dipotassium tetrakis(fluorosulfato)argentate(II) is prepared by the reaction of a 2:1 stoichiometric mixture of potassium fluorosulfate and. silver metal in the presence of the oxidizing mixture, S2°gF2 an<^ HSO^F. In a typical preparation, 0.300 g of silver powder and 0.768 g (5.56 mmol) of KSO-jF were mixed in a one-part pyrex reactor inside the drybox. Approximately 5 ml of each of HSO^F and S20gF2 were distilled in vacuo onto the mixture. Warming the reaction mixture to room temperature and subsequent stirring overnight resulted in the formation of two liquid layers. The clear

S»OcF_ top layer did not seem to mix with the black solid Z o Z

filled saturated HSO^F solution at the bottom. Removal of all volatile materials in vacuum at 50 °C yielded 1.643 g of a black hygroscopic solid powder. Elemental analysis confirmed

the composition as K^AgtSO^F)^. Calculated: %Ag, 18.52;

%K, 13.43; %F, 13.05. Found: %Ag, 18.79; %K, 13.37; %F, 12.90

Alternatively, R"2 [Ag (S03F) ^] was prepared by reacting

Ag(S03F)2 and KSO^F in a one to two mole ratio in BrSO-jF.

Samples of Ag(S03F)2 and KS03F, 0.587 g (1.92 mmol) and 0.530

respectively were mixed inside the drybox as above. Excess

BrS03F was then vacuum distilled onto the mixture. Dissolution was complete after warming the mixture at 5 0 °C for one hour.

The excess BrS03F was then removed in vacuo with the reactor

kept at 50 °C to give 1.114 g of a solid product. This black

solid was identified by its infrared spectrum and melting point

to be K9Ag(SO^F).. 135

3. Experimental Results and Discussions

(a) Infrared Spectra

Much like Ag(S03F)2, the high reactivity of Ag3(S03F)4

and K2Ag(S03F)4 did not allow the use of more conventional

IR mulling agents and window materials. Their black color prevented recording of any Raman spectrum. IR spectra were

obtained only on neat thin powder films between BaF2 plates.

The observed band maxima are listed in Table 20.

The spectra given by both compounds are very similar but also very complex. In the S—O stretching region for a

fluorosulfate group between 1400 and 1000 cm 1, at least seven bands are distinguishable while a single type of fluoro•

sulfate should have a maximum of three bands only. The bands are also sufficiently broad and asymmetric that some fine

splitting may still be present. Two sets of doublets are observed between 900 and 8 00 cm 1 in the S—F stretching region.:

The splittings of the doublets are rather small. Due to the mentioned complexity in the S—O stretching region, detailed assignment and structural conclusion could not be proposed.

(b) Magnetic Susceptibility Measurements

The magnetic susceptibilities of Ag^SO^F)^ and K-jAgCSO^F)^ were measured between ^ 330 and 77 K. Both compounds were found 136

TABLE 2 0

INFRARED SPECTRA OF Ag3(S03F)4 AND K2Ag(S03F)4

1 1 Ag3(S03F)4 (cm ) K2Ag(S03F)4 (cm )

1325 s,sh 1330 s,sh

1278 s 1290 vs

1235 s 1240 s,sh

1195 vs 1190 vs

109 5 mw 1090 w,sh

1082 vs 1080 s

1055 s 1050 vs

855 m 855 m

845 m 845 m

825 s 825

810 s 810

Between BaF plates to be magnetically concentrated. Strong antiferromagnetic coupling is evident from the Neel temperatures of 24 0 and

^ 300 K respectively. The magnetic moments are temperature

dependent and fall below the spin-only value of 1.7 3 yD for one unpaired electron. The results of the measurements are cor

tabulated in Table 21 and a xM versus T plot for Ag3(S03F)4 is shown in Figure 13. Antiferromagnetic coupling is very common in Cu(II) 91 compounds, in particular in carboxylates , but for silver(II) the only example previously reported are the silver(II) bisnico 189 144 145 tinate ' and silver(II) bis(pyrazine) peroxydisulfate

In constrast, the fluoro-analogue of I^AgCSO^F)^ I^AgF^ 154,15

follows the Curie-Weiss law between 2 93 and 82 K, with a magnetic

moment of 1.87 yD at room temperature.

Antiferromagnetism may involve direct metal-metal inter• action or a superexchange process whereby the exchange, proceed through intervening nonmagnetic atoms. This process usually occurs in compounds with small monoatomic anions, X, such as oxides and fluorides, resulting in linear or near linear metal-

X-metal groupings. In addition, the metal must possess half

filled d orbitals of suitable orientation. In view of the 9 bulky size of the fluorosulfate group and the d configuration 2+

of Ag , such intermolecular interaction is rather unlikely.

Direct interactions between magnetic centers may be found in 138

TABLE 21

MAGNETIC SUSCEPTIBILITIES AND MAGNETIC MOMENTS

FOR

Ag3(S03F)4 and K2Ag(S03F) 4

Ag3(SO F) K2Ag(S03F) 3 4 4

cor cor T XM yeff T XM yeff

6 3 1 3 1 (K) (10 cm mol" ) (uB) (K) (10^ cm mol )

309 1130±6 1. 67 ±.02 336 6 8 3 ±7 6 1 .3 5 ±.15

284 1151 ±11 1. 62 ±. 04 307 692 ±59 1. 30±.ll

256 1168 ±9 1 . 55 ±.03 280 689 ±36 1 . 24 ±. 07

232 1168 ±16 1.47±.03 256 683 ±76 1. 18±.13

206 1160 ±8 1.38 ±.01 231 676 ±26 1. 12±. 04

181 1136 ±11 1.28 ±.02 205 660 ±147 1. 04±.23

156 1092±20 1.17±.03 181 632±112 0. 96±.17

131 1022 ±15 1.04±.02 156 592±41 0. 86±. 06

114 966±19 0.94 ±.02 131 534±37 0 .75± . 05

80 939±24 0.77±.02 108 445±15 0 .62± . 02

77 471±32 0. 54 ±.04 FIGURE 13

Temp,[°K] bi- or poly-nuclear complexes.

Assuming a formal dipolar coupling between the metal centers, it is possible to determine the number of magnetic centers over which exchange occurs, n, and the exchange integral, J, by a graphical method. Such a method has been described and illustrated in detail for a number of trivalent 19 chromium and iron carboxylates by Earnshaw, Figgis and Lewis For the S = 1/2 system, the magnetic susceptibility of a number 91 of complexes of bivalent copper as a function of temperature has also been accounted accurately.

The values of \*e£f as a function of temperature (kT/J) for a linear chain of interacting spins of value S, ranging 193 from 1/2 to 5/2, have been calculated for chains of up to

10 members. The listed magnetic moments may be converted to non-ispin-only moments by multiplying by g/2, where g is the tensor determined from the ESR spectrum. For a particular chain length n, comparing the plots of experimental magnetic moment versus kT/J for a range of J values against the theore- 193 tical plot of the tabulated values corrected by the g factor, the best fitting chain length, n, and exchange integral, J, can be determined.

In this study, J values were obtained from the experimen• tal magnetic moments by interpolation using the theoretical plots of calculated magnetic moments versus kT/J for different chain length n. The plot in which best agreement is obtained

for the J values indicated the correct chain length. The average of the interpolated J values from the plot corresponds to the best value for the exchange integral, J. One assumption made in this method is the absolute correctness of the experimentally determined magnetic moments, as every data point is ultilized in the determination of J. If the average

J is within the experimental uncertainties of individual J values obtained, the validity of the method is well confirmed.

If the agreement is unacceptable, the use of the method, and hence the dipolar coupling approach to explain the antiferromagnetism, is probably inappropriate. However, good agreement is very often obtained except at the lower end of the tempera• ture range (near 80 K). This is often explained by the presence of a small percentage of magnetically dilute impuri- 194,195 ties cal

The V ££ versus kT/J plot for n = 2 is illustrated in

Figure 14 for Ag-^SO^F^ as an example. Although the best agreement is obtained from the n = 2 plot, the J values for

Ag^SO^F^ obtained from such a plot did not agree within experimental uncertainties, as shown in Table 22.

In the case of I^AgtSO^F)^, the best agreement was obtainec in the n = 4 plot, the resulting data is shown in Table 23.

Except for the last point, the average J value of - 234 cm-1 FIGURE 14 142

-3.0 -2.0 -1.5 -1.0 -0.5 143

TABLE 2 2

EXPERIMENTAL J VALUES OF Ag3(S03F)4

n = 2

expt. -1, yfiff kT/J J(cm •"•) T

(yB) (K)

1.67 ± .02 -3.00 ± .20 -71.5 ± 5.1 309

1.62 ± .04 -2.41 ± .42 -81.7 ± 17.2 284

1.55 ± .03 -1.97 ± .14 -90.4 ± 6.9 256

1.47 ± .03 -1.64 ± .10 -98.2 ± 6.4 232

1.38 ± .01 -1.40 ± .02 -102.5 ± 1.5 206

1.28 ± .02 -1.21 ± .03 -103.8 ± 2.6 181

1.17 ± .03 -1.05 ± .04 -103.4 ± 4.1 156

1.04 ± .02 -0.89 ± .02 -102.2 ± 2.3 131

0.94 ± .02 -0.80 ± .02 -99.5 ± 2.6 114

0.77 ± .02 -0.67 ± .02 -82.8 ± 2.5 80

J = -93.6 144

TABLE 2 3

EXPERIMENTAL J VALUES OF K2Ag(S03F)4

n = 4

expt.

y -kT/J J(cm 1) -kT/J eff T (y B) (K)

1.35 ± .15 336 1.11 ± .31 -223 ± 86 1.06

1.30 ± .11 307 98 ± .22 -217 ± 63 .91

1.24 ± .07 280 85 ± .12 -229 ± 38 .83

1.18 ± .13 256 75 ± .17 -237 ± 70 .76

1.12 ± .04 231 .66 ± .06 -243 ± 24 .68

1.04 ± .23 205 .57 ± .16 -251 ± 98 . 61

0.96 ± .17 181 50 ± .10 -252 ± 63 . 54

0.86 ± .06 156 .43 ± .03 -251 ± 19 .46

0.75 ± .05 131 37 ± .02 -246 ± 16 .39

0.62 ± .02 108 32 ± .01 -234 ± 8 . 32

0.54 ± .04 77 28 ± .02 -193 ± 15 .23

J = -234 145

is acceptable. This deviation may be caused by very small amount of magnetically dilute impurities. The general better

agreement in the data points within experimental uncertainties

is partially due to the higher uncertainties of individual

data point of the weakly paramagnetic sample.

Using the average exchange integral J for K2Ag(SC>3F)4,

the experimental magnetic moments are plotted against kT/J on

top of the theoretical plot obtained from the tabulated

19 3

values , as shown in Figure 15. Although the deviation

from the theoretical plot is small and within the error limits

except at the lower temperature limit of measurement, the

general curvature and contour of the two plots do not correspond

with each other. Hence it appears that no definite conclusion

can be obtained on the details of the antiferromagnetism of

Ag3(S03F)4 and K2Ag(S03F)4 through the above prodedure.

No ESR signal was observed at room temperature for the

complexes Ag3(SC>3F)4 and K2Ag(S03F)4. At liquid nitrogen

temperature, only broad single line spectra were recorded. The

g-values obtained were 2.140 and 2.173 for Ag3(S03F)4 and

K2Ag(SC>3F)4 respectively. Similarly, the anti ferromagnetic

silver(II) bisnicotinate complex also gives a broad resonance 144 centered at g = 2.08 in the ESR spectrum FIGURE 15 146 147

(d) Oxidation State Determination and Thermal Decomposition

of Ag3(S03P)4

Using the same method as employed for Ag(S03F)2, 1.140 g

(1.584 mmol) of Ag3(S03F)4 was found to give 0.338 mmole of 2-

and consumed 0.1182 mmole of S203 in the iodometric titration. 2 + The amount of Ag present in the sample would be 1.52 mmole and the average oxidation state per mole of silver in the

sample would be 1.32.

Thermal decomposition of 1.74 g (2.42 mmol) of Ag3(S03F)4

under vacuum at 210 °C gave a brownish gas which condensed to

a white solid by cooling with liquid nitrogen. The gas phase

IR identified the 0.20 g of volatile material as So0,F„ J 2 6 2

(1.01 mmol). If the thermal decomposition can be represented

by:

210°C ,

Ag2 [Ag(S03F)4] 3 Ag(S03F) + 2 S206F2 (3.15)

The expected amount of S20gF2 released would be 0.24 g (1.21 mmol)

while 1.50 g of Ag(S03F) would be the non-volatile residue. The actual greyish solid residue weighed 1.54 g (7.44 mmol) and

was identified to be a rather impure AgS03F by its infrared

spectrum.

Trisilver tetrakisfluorosulfate is thermally stable up

to 170 °C under one atmosphere of dry nitrogen and decomposes to give S20gF2 and AgSO^F on further heating. The thermal

stability of K2Ag(S03F)4 is slightly higher than that of

Ag(S03F)4 and did not melt with decomposition below 195 °C.

Formation of S2(-,gF2 was again detected by the gas evolution and by the gas phase IR spectrum.

E. SILVER(II) HEXAKIS(FLUOROSULFATO)METALLATE(IV)

1. Introduction

The analogy between fluoride and fluorosulfate derivatives of divalent silver can be extended to cationic complexes. With

a dipositive cation in Ag(SC>3F)2, donor-acceptor type complex formation with a tetravalent metal fluorosulfate appears feasible. As numerous examples of Group IV A, B and Group VIII

IV II IV metals are known to form M2(M Xg) or M (M Xg) complexes.

In particular, the analogous ternary fluoride complexes of silver(II) have been synthesized as the complexes AgII(MIVF,) 6 where MIV = Ge, Sn, Pb, Ti, Zr, Hf, Pd, Pt 151. Furthermore, 42 the hexakisfluorosulfato complexes of tin(IV) and more 196 recently platinum (IV) such as K_ [Sn (SO..F) c ] and Ba [Pt (SO_,F) , 2. JO j b have been synthesized. Their vibrational spectra should help in characterizing the corresponding silver complexes. In 149

119 addition, Sn Mossbauer Spectroscopy is available for studying the tin complex. Hence the silver(II) fluorosulfate complexes of platinum(IV) and tin(IV) were synthesized and described in the following section.

2. Syntheses and Elemental Analyses

(a) AgPt(S03F)6 71

A solution of Pt(SC>3F)4 in HSO^F was conveniently obtained by the oxidation of platinum metal powder by S_0,F_ 2 6 2

196 in HS03F at 100 °C . such a solution prepared from 0.253 g of platinum powder in a one-piece pyrex reactor was then added

in the drybox to an equimolar suspension of Ag(S03F)2 in a

1:1 by volume mixture of S~GvF_ and HSO..F. The Ag(SO-,F)„ z o / 3 32 suspension was prepared separately from the reaction of 0.140 g of silver powder with a S~0,F_ solution in HS0oF. z 6 z 3 The resulting mixture was stirred at room temperature for one day. Subsequently S„0,F„ was first distilled off in z b z vacuo, then the light green precipitate was isolated by vacuum filtration. Remaining traces of volatile materials were removed in vacuo. A homogeneous light green powder [1.050 g,

1.194 mmol of AgPt(S03F)g] was obtained. Elemental analysis,

calculated for AgPt(S03F)6: %Ag, 12.02; %Pt, 21.74; %F, 12.70.

Found: %Ag, 12.30; %Pt, 21.52; %F, 12.89. 150

Silver(II) hexakis(fluorosulfato)platinate(IV) was found to be very hygroscopic and thermally stable to 110 °C.

Between 110 and 180 C, the material turned reversibly brown and decomposed on further heating to a red orange solid.

Alternative synthetic methods such as oxidizing a stoi• chiometric mixture of silver and platinum metal; addition of silver metal to a solution of Pt(SO-.F). in a mixture of S„0,F„ 3 4 2 6 2 and HSO^F or platinum metal to a suspension of Aq(SO^F)^ in

^2°6F2~~HS°3F' ^"""^ not result in a homogeneous product. Dark brown or black particles were mixed-in with the light green product. These were most likely due to coating of metal surfaces that prevented further reactions.

(b) AgSn(S03F)6

Metallic tin (0.257 g) was reacted with about 10 ml of a

2:1 mixture by volume of HS0oF and S„OrF„ in a one-part pyrex

j Z 6 Z reactor at room temperature for one day. To the resulting milky white suspension, the stochiometric amount, 0.233 g of silver powder was added inside the drybox. The mixture was stirred at room temperature overnight. A homogeneous green solid was formed. Removal of all volatile material in vacuum yielded 1.773 g of a hygroscopic green powder compared to 1.776 g expected for AgSn(SO^F)^. Elemental analysis, calculated:

%Ag, 13.14; %Sn, 14.46; %F, 13.89. Found: %Ag, 14.27;

%Sn, 14.56; %F, 14.03. 151

Silver(II) hexakis(fluorosulfato)stannate(IV) was found to melt above 170 °C with decomposition to a white solid and brown gas which condensed to a colorless liquid, judged to

be SnO,F_. 2. 6 2

3. Characterizations

(a) Vibrational Spectra

The relatively light green color of the platinum and tin complexes synthesized above allowed the recording of excellent

Raman spectra. The Raman spectrum of AgPt(S03F)g is shown in Figure 16 as an example. In addition, infrared spectra of

neat sample powders between BaF2 and KRS-5 plates were obtained.

The vibrational band maxima of AgPt (S0oF) ,_ and AgSn(S0oF), J b 5 b are listed in Table 24 along with those of some related compounds.

Assignment for the vibrational bands becomes very difficult for the complexes of the type Ag^M"1^ (SO..F) , where MIV = Pt J b or Sn. Only an approximate description of the SO^F group vibra• tion is possible, because proliferation of bands due to slight non-equivalence or vibrational coupling is expected. The overall band shapes for the two complexes are very similar.

When the spectra are compared to those of analogous complexes, 196 4 2 BaPt (SO-.F) and K„Sn(SO-F), , as shown in Table 24, bands -J b z 5 b at ^ 1400, ^ 1250 and ^ 1000 cm 1 can be interpreted as being due to monodentate OSO^F groups linked to platinum or tin; 152

FIGURE 16

The Raman Spectrum

of Ag Pt(S03F)6 from 100 to 1500 cm-' TABLE 24 VIBRATIONAL SPECTRA (cm ) OF AgPt (SO-.F) _ , AgSn (SO-.F) , , AND RELATED COMPOUNDS H

A Pt(S F) A 5n(S F) g °3 6 Ba[Pt(S03F)6] g °3 6 K2[Sn(SO3F) ]

b „ T„c T„d Raman IR KRS-5 plates Raman Raman IR IR

1406 m 1410 vs 1410 vw 1412 ms 1425 s, sh 1380 s, b 1380 m 1380 s, sh 1397 msh 1404 s 1395 vs , b 1220 s, b 1249 vs 1215 vs, b 1386 s 1398 s, sh 1227 s 1190 s, b VL220 m, sh 1148 vs, b 1258 vs 1263 mw 1170 vs , b 990 s, b 1209 mw 1080 w 1218 ms 1192 w 1040 s 820 s, sh 1125 m 1030 w, sh 1033 s 1120 vs 1010 s, sh 810 s, b 1048 ms 968 vs, b 1012 ms 1013 ms 984 s, b 620 s 1017 mw 925 s %950 vw 9 8 0 vw, sh 865 s, sh 571 m 962 w 850 m, sh 857 ms 850 m 825 s, b 550 s ^930 vw 830 s 838 w 820 w, sh 640 s, b 430 m 850 w 760 vw 629 vs 633 ms, sh 625 s, sh ^820 w 738 vw 583 w 620 ms 590 s 629 vs 660 s 549 m 590 s 555 s 589 mw 630 mw 460 s 560 ms 452 m, sh 549 mw 590 s 422 w, sh 523 mw 448 s 520 mw 550 s 411 ms 434 s 443 s 520 vw, sh 283 s 413 ms 480 w 468 m, sh 210 ms 278 ms 273 s 452 s 178 m 212 s 305 s 158 m 295 m, sh 146 mw 135 mw

3. be See Appendix A for abbreviations. Reference 196. Composite spectrum obtained on solids

between B'aF2 and KRS-5 windows. ^Reference 42. 154

additional S—0 stretching vibrations at ^ 1150 and ^ 1040 cm-1 are observed for the silver(II) complexes. Bands in this region are best assigned to bidentate fluorosulfate groups, and the reason for their appearance must be seen in the strong 2 + polarizing effect of the Ag ion.

The occurrence of both mono-and bidentate SO^F groups in the silver(II) complexes indicates that not all the fluoro• sulfate groups are coordinated to the MIV ion; some appear II IV bonded to both Ag and M . In any event, these two silver(II) complexes are best regarded as ternary fluorosulfates with the fluorosulfate groups coordinated to both metals rather than ionic complexes of the type Ag11 [MIV (SO-.F)such as the 3 b example of K„ [Sn (SO-.F)• , ] . 2 6 b (b) Electronic Spectra

The diffuse reflectance and mull spectra were obtained for the complexes AgPt(SO-.F) and AgSn(SO-F),. Resolution in i b i b general was rather poor. Curve-resolving techniques have been used to extract ligand field parameters from the fluoro- 152 153 analogues of these complexes ' , resulting in a 10 Dq value of 11,700 cm for AgSnF^. However, where only a broad asymmetrical band is encountered, as in the case of AgSn(SO-,F)^ 3 6 shown in Figure 17, rather than the presence of shoulders, data obtained by curve fitting methods may not be very reliable 155

FIGURE 17

350 400 450 500 550 600 650 700 750 Wavelength [nm] 156

Although some similarity is found for the electronic spectra shown in Table 25, unlike AgSnF^, no assignment can be

readily made to the broad bands of AgPt (SO-,F) , and AgSn (S0oF) , . J b J b

But in comparison to the spectrum of AgCSO^F^, the d—d tran•

sitions seem to occur at lower wavelengths, with the band center shifted from 22,000 to ^ 16,000 cm ^. Weakening of the ligand field around silver when the SO^F groups are coordinated 2+ 4+ 2+ to both Ag and Sn or Pt may be responsible for such a reduction in ligand field splitting.

Both AgPt(S03F)g and AgSnCSO^F)^ were found to obey Curie-

Weiss law between ^ 300 and ^ 80 K with Weiss constants of - 0.7 cor and - 5.8 K respectively. The plot of l/xM versus temperature

is illustrated in Figure 18 while the magnetic susceptibility data is listed in Table 26.

Substantially higher Ve^^ values are found for the two complexes, where the room temperature magnetic moments are

1.94 and 2.18 y„ for AgSn(SO-)F), and'AgPt(SO,F), respectively. B j b J b

Although BaPtCSO^F)^ is diamagnetic consistant with d^ spin-paired configuration, small contributions from PtIV to the paramagnetism of AgPt(SO^F)^ cannot be ruled out entirely.

Such contribution could result from temperature independent 157

TABLE 25

ELECTRONIC SPECTRA OF AgPt(SO,F),, AgSn(S0oF)^ AND RELATED COMPOUNDS J b 5 b

-1 3 Compound Type of Spectrum ^max ^cm x 10 ) Reference

AgPt(S03F)g Diff. refl. ^25.6, 16.0 This work

AgPt(S03F)6 Mull ^25.0 This work

AgSn(SO F) Mull M 0 . 0 , 22.5 sh This work 3 16.0, 12.5

AgSn(SOnF)^ Diff. refl. >28.5, 16.1 This work 3 6

AgSnF, Diff. refl. 15.3, 13.7 153 6 11.7, 8.05 158

FIGURE 18

Magnetic susceptibilities of

600, AgPt(SQ.F)cond AgSn(SQ&

500

'AgSn(Sg3F)6 C = 0.469+ 0.004

4004 8= -5.8+ \B°K J /i.eff= 1.94+0.02 B.M corr X m 300H

200-

oAqR(S03F)6

100 A C = 0597 + 0.005 B = -0.7+ i.6°K extrapolated

/ /*ff= 2.19+0.02 B.M.

~\—r 0 100 200 300 159

TABLE 2 6

MAGNETIC SUSCEPTIBILITIES AND MAGNETIC MOMENTS

AgPt(S03F)6 and AgSn(S03F)6

AgPt(S03F) AgSn(S03F) 6

cor a cor a T XM yeff T XM ueff (10^ cm mol ^) 1 (K)

301 1994 2.19 301 1539 1. 94

275 2168 2.19 275 1676 1 .94

249 2381 2.18 250 1827 1.93

225 2633 2.18 225 2026 1.94

200 2980 2.18 203 2256 1 .94

176 3347 2.18 175 2573 1 .93

151 3895 2.17 153 2939 1.93

128 4570 2.17 152 2943 1.93

107 5633 2. 20 128 3441 1.92

80 7602 2. 21 109 4187 1 .96

79 5704 1.96

aMagnetic moments are calculated by using the Curie-Weiss law:

r 1/2 ueff = 2.828[X^° (T - 0)] paramagnetism (T.I.P.). Similarly, the ternary fluorides of

IV the type AgM Fg with M = Sn, Pb, Pd are magnetically dilute

and show room temperature moments in the range of 1.92 -

1 1 1.99 yD ^ , but for AgTiFc, a rather high temperature inde- 151 pendent value of ^ 2.21 y_a is found between 296 and 100 K

Already mentioned when discussing the electronic spectra,

both electronic delocalization and a decrease in 10 Dq, as

indicated by shifts in d—d bands, seem to cause an increase

in yeff in the series [Ag (bipy)2 ] (S03F)2 , Ag(SC>3F)2 and

IV IV AgM (S0oF), with M = Sn, Pt. 2,2'-bipyridine is expected to o b

create a much stronger ligand field than S03F , but the differ•

ence displayed by the fluorosulfate group depending on its 2+ 2+ 4 + 4 + coordination to Ag only or to both Ag and Sn or Pt is worth noting.

(d) ESR Spectra

Like the corresponding ternary fluorides, the two fluoro•

sulfate complexes AgPt (SO-.F) and AgSn (SO-.F) , gave anisotropic J b j b

spectra with clearly resolved g|| and g^ components at both

295 and 80 K. The g-values along with the calculated magnetic moments using equation (3.9) are listed in Table 27. Again good agreement between calculated magnetic moments and the

experimental values is found except for AgPt (SO-.F) where J b bulk magnetic susceptibilities indicated a higher magnetic TABLE 27

ESR DATA OF AgPt(S03F)6, AgSn(S03F)6 AND RELATED COMPOUNDS

Compound T,K y V Ref. eff, B

AgPt(S03F)6 (s) 295 2.266 2.494 2.143 1 . 96 This work

AgPt(S03F)6 (s) 80 2. 258 2.486 2.134 1 .96 This work

AgSn(S03F)6 (s) 295 2.255 2.481 2.134 1.95 This work

AgSn(S03F)6 (s) 80 2. 245 2.480 2.117 1.94 This work

AgSnF, (s) 80 2.315 2.610 > 2.153 2. 00 152 b

AgHfFg (s) 80 2. 275 2. 519 2.143 1.97 152

a 1/2 Calculated from the expression: ^eff = g [S(S + 1)] moment. The discrepancy is consistent with small paramagnetic contributions from the platinum containing moiety.

ii: (e) 'Sn Mossbauer Spectrum of AgSn(S0oF) 3 6

As the spacings of the nuclear energy levels depend on the electron distribution about the nucleus, especially the valence electrons, the energy difference between the ground and the first excited state of the source and absorber in a

Mossbauer experiment will be different. Resonant reabsorption by the absorber of the y ray emitted by the source can be achieved by moving the source relative to the absorber to modulate the frequency of the y emission by the doppler effect

The isomer shift, 6, is the measure of the source velocity needed to bring the sample absorber into resonance with the source and depends on the radius of the nucleus and the s elec• tron density at the nucleus'. An increase in 6 corresponds to an increase in s electron density at the absorbing tin nucleus

Variations in isomer shift are reported relative to SnC^, customarily used as the reference and set at 0 mm/sec. 119

The first excited nuclear state of Sn has a nuclear moment, I = 3/2, so that tin compounds will have a quadrupole splitting, A, whenever an imbalance of charge around the tin nucleus causes an electric field gradient, q. The size of q depends mainly on the imbalance in the distribution of the 163

valence electrons on tin, [q ,], which is caused by the va i different orientation of ligands around tin and the imbalance

in the polarity of the tin-logand a bonds. The a bond polari- * 44

ties can be estimated by the Taft inductive constants, a , mentioned in Section IB. To a much lesser extent, q also

depends on the charges on the atoms surrounding tin, [q, .], lat which has a small cancelling effect on [q n]. For compounds va i with the same geometry, the bond polarity differences between the ligands increases with A. 119 The Sn Mossbauer spectrum of AgSn (SO-.F) at 80 K shows J b a quadrupole splitting of 0.5 mm/sec. Previously studied

compounds containing the Sn(S0oF), group as in M^[Sn(SO-F),] J b 2 j b I + • + + +42 + +197 where M = C102 , K , Cs , NO or I(S03F)2 , Br(S03F)2 ,

all gave single line Mossbauer spectra. The observed quadrupole

splitting for AgPt(S03F)g is less than that of the polymeric

Sn(S03F)4 where a value of 1.34 mm/sec is found.

The isomer shift for AgSn(S03F)g is - 0.211 mm/sec relative

to Sn02, slightly higher than previously reported values of

1 42 197 between - 0.23 and - 0.30 mm/sec for M [Sn(SO_F) ] compounds ' / 2 3 b but the difference is probably not significant in view of the accuracy limit of ± 0.03 mm/sec.

The small quadrupole splitting for AgSn(S0oF). most 3 b likely results from the unequal contribution of the six fluoro•

sulfate groups around each tin atom, in contrast to the regular 2- SnO, coordination octahedral in the other Sn(SO-F),, b 3 b 4 2 192 species ' . This difference in interaction between differ ent fluorosulfate groups is caused by the bonding of some

SO^F groups to tin(IV) only with others sharing between both tin(IV) and the strongly polarizing silver(II) ion, thus weakening the contribution to the tin(IV) nucleus. Such a result is consistent with the observed trend of increasing

IV magnetic moments going from [Ag(bipy)2] (SO^F)2 and AgM (S03F) as noted previously.

F. REACTIONS OF Ag(S03F)2

1. Reaction of Ag(S03F)2 With Elemental Fluorine

(a) Introduction

Silver(II) fluoride acts as a catalyst in the fluori• nation of sulfur trioxide to give bisfluorosulfuryl peroxide,

S F an< 2°6 2 ^ fluorine fluorosulfate, FOS02F according to:

+ 160 °C 164 2 S0o + F„ ^ S„OrF„ (3.16) 3 z , _ z b z AgF0

+ 220 °C fi3 and SO + F *• FOSO F lbJ (3.17)

AgF2 165

Even though the detailed mechanism of such a catalytic fluorination is not known, the intermediate formation of

AqiSO^F)^ or perhaps FAgSO^F is likely. The successful

insertion of SO^ into AgF2 to give Ag(SC>3F)2 described in

Section B2(c) of this chapter is consistent with this view.

It seemed therefore appropriate to carry out the reaction

of Ag(S03F)2 with elemental fluorine to further investigate such a proposed mechanism.

(b) Conversion of AgtSO^F^ into AgF2

A two-part monel metal reactor (Section II) containing

0.8 g of Ag(SC»3F)2 was filled with reagent grade fluorine to atmospheric pressure at 25 °C. Also connected to the metal vacuum line was an evacuated one part pyrex trap cooled under liquid nitrogen. The progress of the reaction was monitored by weighing the metal reactor after first removing the fluorine by pumping on the reactor cooled by liquid nitrogen and then condensing the less volatile materials into the pyrex trap.

No detectable reaction occured after heating the reaction mixture at 55 °C for one hour. Raising the reaction temperature to 100 °C for one hour gave no noticable reaction. However, after heating the reaction mixture at 130 °C for two hours and removal of all excess fluorine, volatile materials condensed

s F and in the pyrex trap as white solid. Both 2°g 2 FOS02F 166

along with traces of SiF^ were identified by the gas phase IR spectrum of the volatile material. The solid residue had no

IR bands due to SO^F groups, in fact, no absorption was detected

down to 800cm , the limit of transparency of the BaF2 plates used. This dark brown-black residue showed its oxidizing ability by oxidizing iodide to iodine. On hydrolysis, no sulfate ion was detected when tested with aqueous barium nitrate. The silver analysis indicates the residue to be AgF,,:

calculated , %Ag, 73.95 for AgF2; found: %Ag, 74.09.

The result of the above experiment and the insertion of

SO^ into AgF2, represented by equation (3.18), (3.19) and

(3.20), favours Ag(S03F)2 as an intermediate in the synthesis

S F and of 2°6 2 FSO^F shown by equations (3.16) and (3.17).

+ 50 °C

AgF2 + 2 S03 *- Ag(S03F)2 (3.18)

+ 130 °C

Ag(S03F)2 + F2 — AgF2 + S2OgF2 (3.19)

+ 130 °C

Ag(S03F)2 + 2 F2 AgF2 + 2 FOS02F (3.20)

Even though the actual synthetic reactions (Equation 3.16 and 3.17) are carried out at higher temperature, it seems reasonably to assume faster conversion rates between Ag(S03F)2

and AgF2 at higher temperature, based on the observed pyrolysis

of Ag(S03F)2 to give S2OgF2 and also the direct fluorination

of metallic silver and silver(I) salts to give silver(II)

difluoride mentioned in Section III.A.

2. Reactions With Some Other Halogens

(a) Chlorine

Excess dry chlorine ( ^ lg ) standing over ^2^5' was

distilled onto 0.189 g (0.617 mmol) of Ag(S03F)2 in a one- piece thick wall pyrex reactor. On warming to room temperature

the dark brown Ag(S03F)2 turned gradually to black, then

greyish-white within four hours. The low temperature IR

spectrum of the volatile material at 77 K showed bands due to 166

C10S02F . After all volatiles were removed, 0.133 g of

I white solid was left behind. The expected weight for Ag (S03F) would be 0.128 g (0.617 mmol). The IR spectrum of a nujol mull sample of the white solid is identical to that of

I Ag (S03F) listed in Table 16. Hence, + 25 °C

2 Ag(S03F)2 + Cl2 (excess) 2 ClOSO^F + 2 AgSO-jF . . . . (3

(b) Bromine

In much the same manner as in the reaction with Cl excess dry Br2 was vacuum distilled onto AgCSO^F)^ in a one part pyrex reactor. On warming to room temperature, the

mixture reacted, turning the Ag(S03F)2 to a non-homogeneous

black and white mixture. Further reaction with more Br2 resulted in a white precipitate in the red suspension. Before all the volatile materials could be removed by vacuum pumping, part of the white material turned black again. A low tempera• ture IR spectrum of the volatiles identified the less volatile

198 component to be BrSO^F

It appears that Ag(S03F)2 oxidized Br2 to BrSO^F, initiall

yielding AgSO^F. Due to its higher volatility, the excess Br2 was preferentially removed by vacuum pumping, leaving behind the less volatile product BrSO^F, which reoxidized part of the

AgSO^F to some Ag(II) species, presumerably Ag^CSO^F)^.

A similar reaction was carried out with Ag^(SO^F)4.and

excess Br2. Identical result was obtained. No homogeneous solid product could be isolated due to the above mentioned

difference in volatility between Br2 and BrSO^F.

This system was further investigated by reacting Ag(S03F)2

with a mixture of excess BrSO^F and Br2 ( ^ 1:1 by volume).

Initial reaction was similar to that of Ag(S03F)2 and Br2, a mixture of black and white precipitate appeared. But when the volatile materials were slowly distilled off while the mixture was being stirred, a homogeneous black precipitate remained. 169

This black solid was identified to be Ag^CSO^F)^ by its IR spectrum and thermal decomposition.

Hence the result from the above reactions provided an

interconnecting relationship between silver metal, Ag(SC>3F)2,

Ag^CSO^F)^ and AgSO^F, connected by the reagents Br2, BrS03F and S„GvF„. Such a relationship is illustrated in Figure 19. 2 o 2

3. Reactions With Chloryl Fluorosulfate

To investigate the ability of Ag(S03F)2 as a fluorosulfate ion acceptor further and perhaps to help clearify the structure

of the anionic complexes Ag2 [Ag (SC>3F)4 ] and K2 [Ag (SC>3F) ^] ,

C1C>2S03F was ultilized as a donor. C102S03F readily donates

a fluorosulfate ion to a number of acceptors such as Sn(SC>3F)4 to form the complex (C1C> ) Sn (SC> F) , containing the heterocation 2 3 6

Due to the low volatility of C102S03F in vacuo, it was

pipetted ( ^ 2 g ) onto 0.640 g (2.01 mmol) of Ag(SC>3F)2 in a one part pyrex reactor inside the drybox. A solution was obtained after the resulting mixture was stirred at room temperature for one day. Volatiles were then removed by pumping under vacuum at room temperature for one day. A black solid (1.015 g) remained as compared to 0.989 g expected for

ClC>2Ag (SC>3F)3 . The sample appeared slightly wet between the 170

FIGURE 19

INTERCONNECTING REACTION SCHEME

BETWEEN THE FLUOROSULFATES OF SILVER 171

BaF^ windows and contained both bidentate and monodentate

+ fluorosulfate groups in addition to the bands due to the C1C>2 198

cation . Further pumping on the black solid at room tempera•

ture again reduced the weight of the sample but dark brown

impurities started to appear. Over a period of one week, the

sample weight reduced to the original value of the starting

material while its physical appearance corresponded to the

dark brown Ag(S03F)2. Hence it appears that Ag(S03F)2 is not

a strong SO^F ion acceptor and gradually loses ClO^SO^F in

vacuum in a reversal of the complex formation reaction.

25 °C

C102S03F + Ag(S03F)2 ^ C102Ag(S03F)3 (3.21)

Another attempt was made by adding S„OrF„ into the mixture 2. b 2

of C102S03F and Ag(S03F)2 to hopefully further oxidize the

silver(II) ion thus forming anionic complexes analogous to

I III 8 Hoppe's M Ag F4 "*"^ . However on removing the volatile mate•

rial, Ag(S03F)2 was again obtained.

4. Reaction With Pyridine

As demonstrated in Section III.C, Ag(S03F)2 can be stablized by nitrogen donor ligands such as 2,2'-bipyridine, it appeared a similar complex could be prepared using pyridine 172

as the ligand and serving as the solvent in the reaction. As

mentioned in Section III.A., many pyridine complexes of

•i lTT^ u K ^ ^ 122,123,125 silver(II) have been reported

When excess dry pyridine ( ^ 10 ml ) was vacuum distilled

onto 0.288 g of Ag(S03F)2, a dark red-brown solution was

obtained after warming to room temperature. Attempts to remove

the excess pyridine in vacuo completely resulted in a non-

homogeneous red-brown and white slush. It appears that the

pyridine was attacked by AgfSO^F)^ and no simple coordination

complex could be isolated.

5. Reaction With Acetonitrile

Acetonitrile has been found useful as the solvent in

synthesizing [Ag(bipy)2] (SO^F),,. Although CH^CN can also serve 2 +

as a ligand, the Ag ions are coordinated by the 2,2'-bipyridine

ligands only. It seemed interesting to investigate the reaction without the bipyridyl ligand.

Dry CH3CN ( ^ 10 ml ) was vacuum distilled onto 0.289 g of AgtSO^F),,. Stirring at room temperature resulted in a dark brown solution with small amounts colorless solid gradually precipitating out. Vacuum removal of volatiles resulted in a very viscous dark brown liquid, which turned into a dark brown sludge on warming at 60 °C while pumping. White impurities began to appear on further heating and pumping. It appears 173

that the solution of Ag(SC>3F)2 in CH3CN is rather unstable

and no complex could be isolated.

6. Reaction With Antimony Pentafluoride

Antimony pentafluoride, an exceptionally strong Lewis

acid, has found the greatest use in the formation of hetero-

cations such as the poly-interhalogen cations, mainly through

its strong acceptor ability via fluoride abstraction. More

recently, SbFj- has been proven to be a good fluorosulfate

199

acceptor, useful in solvolysis reactions of ClC^SO-jF ,

200 200 IC12S03F and IBr2SC>3F in SbF5 to give respectively the

heterocation containing ClO^b^^, ICl-jSb^^ and IBr.jSb.jF-^

complexes, along with some antimony(V) fluoride-fluorosulfate 201

such as Sb2F^S03F resulting from the fluorosulfate abstrac•

tion process. The behaviour of Ag(S03F)2 in SbF^ would be

interesting in view of the possible fluorosulfate abstraction

to form silver(II)-antimony(V) donor-acceptor complexes.

Excess purified SbF^ ( % 10 ml ) was vacuum distilled onto

1.264 g (4.131 mmol) of Ag(S03F)2 in a one piece pyrex reactor.

On warming gradually to room.temperature, the mixture turned

green initially, then continued to turn white slowly when continuously stirred at room temperature overnight. Homogeneous white suspended particles were found in the viscous liquid 174

SbF 5. A small amount of non-condensable gas was detected, assumed to be oxygen, and subsequently pumped off. The presence 202 170 of S02F2 and SiF4 in the volatiles was identified by the gas-phase IR spectrum taken between AgCl windows. The

infrared spectrum of the viscous liquid between BaF2 windows showed bands at 1430, 1125, 1070 and 890 cm"1 in addition to others at 1165 cm 1 and shoulder at 1025 cm-1. The first four

bands were most likely due to Sb2Fg(S03F) when compared to its 201

IR spectrum . Removal of all volatile materials from the reaction mixture resulted in a white powder. No absorption was observed.in the S—O stretching region of the IR spectrum of the white solid, while the broad bands at ^ 660 cm 1, 340 cm 1 and ^ 475 cm 1 were the only bands observed down to 250 cm 1. The first two are best assigned to Sb—F vibrations of a SbF " 203 6 ion

The initial green color suggests perhaps complex formation, reminiscence of the other ternary fluorosulfate complexes of

Ag(II), AgMIV(SO^F) . However such a complex appears to be

AgSbFunstabl6. e and breaks down into the detected oxygen, S02F2 and 175

G. ATTEMPTS TO OBTAIN HIGHER OXIDATION STATES OF SILVER

Having been unsuccessful in obtaining compounds with

silver in an oxidation state higher than two via direct

oxidation by 820^2, it appeared appropriate to attempt the

conversion of the existing silver(III) containing compounds

into fluorosulfates.

S F 1. The reaction of AgO and 2°6 2

The mixed valency silver oxide, Ag[AgIII02], commonly

known as AgO, contains silver(III) and silver(I) centers. Its

reaction with S_OrF„ would hopefully lead to some silver(III) Z b Z containing fluorosulfate or mixed valency silver(II), (III)

systems.

Excess S~OrF„ ( ^ 5 ml ) was distilled into a one-part

Z O Z thick-wall pyrex reactor containing 1.445 g (11.67 mmol) of AgO.

Large amount of uncondensable gas was evolved when the mixture was allowed to warm to room temperature. This was judged to be oxygen and subsequently pumped off. After stirring the mixture overnight and removal of the gas evolved, the mixture was stirred at 50 °C for a further two days until gas evolution ceased as evidenced by no detectable pressure at liquid nitro• gen temperature. After distilling off the excess S^O^-F^, 3. 602 g 176

of dark brown powder remained. This yield is comparable to

the 3.569 g expected for Ag(S03F)2< The product was further identified by its IR spectrum and its melting point.

Although the oxidation of the Ag(I) ions to Ag(II) was predicted, the reduction of Ag(III) to Ag(II) is rather surpris•

ing in such a medium. The detailed mechanism of the reaction is unclear, an initial coproportionation to Ag(II) may be possible.

2. The Reaction Of CsAgF4 And S03

(a) Introduction

As mentioned in Section A of this chapter, silver(III) complex fluorides of the type JYT^Ag113^, with M = K or Cs, has 158 been synthesized and characterized . In view of the success•

ful conversion of AgF2 to Ag(S03F)2 by S03 insertion, as described earlier, the possible formation of a fluorosulfato- argentate(III) complex via the reaction of a tetrafluoro- 158 argentate(III) complex and sulfur trioxide should be investigated. Since silver difluoride is found to be an effective 128 12 9 fluorinating agent ' and also as catalyst in fluorination 130 reactions , CsAgF^ should be an even more effective fluorin• ating agent. 177

(b) Synthetic Reactions

In a typical reaction, approximately 1.8 g (5.7 mmol) of

CsAgF4 (Section II.C.2.) was reacted with an excess (about

3.1 g) of SO.^. Sulfur trioxide was first distilled from oleum

under a dry nitrogen atmosphere, then vacuum distilled into

the reactor, a 150 ml two part monel metal can equipped with a

Hoke valve and a Teflon coated stirring bar. The reaction mixture was stirred at 25 °C for one day. The gas phase infrared

spectrum of the volatile materials showed besides traces of 75

SiF^ and SO^, only bis(fluorosulfuryl) peroxide, S2°6F2 ' as

the fluorinated reaction product.

Removal of the remainder of the volatile materials, including the excess SO^, was accomplished by heating the reactor to 60 °C in vacuo. Approximately 3.1 g of a dark greenish brown powder of the composition CsAgtSO^F)^ (5.7 mmol) were obtained. The material was extremely moisture sensitive and showed thermal stability up to 130 °C (then decomposed to a black solid). The elemental analysis established the composit-

tion as CsAg(SC>3F)3. Calculated: %Cs, 24.71; %Ag, 20.05;

%S, 17.88; and %F, 10.59. Found: %Cs, 24.84; %Ag, 19.97;

%S, 17.68; and %F, 10.37.

The overall reaction may be formulated as:

24 Hr

2 CsAg + 8 S03 (3.23) 178:

It seems that both SO^ insertion and fluorination take place

even at room temperature. A precedent for such a reaction is 8 8 reported by Brown and Gard :

CrF5 + 5 S03 mm Cr(S03F)3 + S2OgF2 (3.24)

Both reactions together with others, like the photolysis of 204 C10S02F and the thermal decomposition of xenon fluorosul•

ZU3205,20zuo6 fate ' / may be regarded as alternative routes to S206F2.

88 204 205 206 A point frequently stressed ' ' ' is that these methods

do not involve the use of elemental fluorine, which is poten•

tially highly hazardous and requires the use of metal1 high

vacuum lines not commonly available in many laboratories.

However this is somewhat misleading, since in all cases the

synthesis of the starting materials will involve elemental

fluorine at certain stages of the synthetic scheme. For example, 207

CrF5 is prepared by the reaction of CrF3 and F2; while ClF, 2 08

used in the preparation of C10S02F , is synthesized from its corresponding elements. Only the anodic oxidation of 209

S03F in HS03F reported by Dudley avoids entirely the use

of elemental fluorine in the overall synthesis. But none of

the alternative routes to S2°6F2 all°ws the convenient, large

scale synthesis via the catalytic fluorination by AgF2 discussed in Section III.F.l. 179

To identify the product as CsAgCSO^F)^, an alternative synthetic route was attempted. A 1:1 stoichiometric mixture of silver metal powder and CsSO^F was reacted with the oxidi•

zing mixture of HSC> F and S„C> F_ to hopefully obtain CsAg(S0oF)_ 3 z b z ~ J 3 according to:

HS03F-S206F2

Ag + CsS03F »• CsAg(S03F)3 (3.25)

However the resulting brown-black solid product appeared non-homogeneous after the volatiles were removed. Instead

of CsAg (S03F) ,• a mixture of Cs2Ag(SC>3F)4 and Ag(SC>3F)2 was probably formed:

HSC> F-S-O,.F_ 3 2 b 2

2 Ag + 2 CsS03F Cs2Ag(S03F)4 + Ag(S03F)2

(3.26)

The high reactivity of CsAg(S03F)3 allowed the recording

of infrared spectrum only on thin solid film between BaF2 plates, while no Raman spectrum could be obtained, most likely due to its dark color. The infrared band positions of

11 CsAg (S03F)3 and those of some related compounds are listed in

Table 28. A tracing of the spectrum is shown in Figure 20.

The observed band positions of CsAg (SO-.F) differ from those of FIGURE 20

The I.R. Spectrum of

Cs CAg(S03F)3] in the BaF| region

i 800 cm-i 181

TABLE 28

1 INFRARED SPECTRA (cm ) OF CsAg(S03F)3 AND RELATED COMPOUNDS

a 8 2 3 CsAg(S03F)3 CsS03F Ag(S03F)2 Assignments

1365 s

1300 s 1320 s, b vS03 (A") 1330 vs

1200 b, vs 1258 s 1185 vs, b vS03 (A')

1080 s, sh

1071 s 1070 s, b vS03 (A') 1050 vs

810 ms 728 s 820 ms VS-F (A')

a. between BaF2 plates 182

Ag(S03F)2. With the absence of any ionic fluorosulfate bands,

82 in particular, those of CsSO^F ,74^ ^e possibility of

CsAgtSO^F)^ being a simple mixture of Ag(S03F)2 and CsSO^F can be ruled out. In addition, a sample exposed briefly to moist air showed an addition band at 1270 cm 1, indicative of ionic fluorosulfates. The number of S—O stretching bands

observed suggests a Cs symmetry for the fluorosulfate groups.

The observed band positions, with allowance for site symmetry splittings, appear to be most consistent with bidentate fluoro- 7 fi

sulfates, as in the case of (CH3)2Sn(S03F)2 (Table 10). This implies a distorted octahedral environment around silver.

The magnetic susceptibility of CsAg(S03F)3 between 300 and 77 K was measured and the result is listed in Table 29.

A magnetic dilute system which shows Curie-Weiss behaviour was observed. The Curie constant C was 0.4 61 ± 0.003 with the

Weiss constant, 0, extrapolated to be 3.7 ± 1.3 K. The observed magnetic moment 1.91 yD at room temperature is com• es parable to other similar magnetically dilute silver(II) systems such as a value of 1.92 y_, for Ag(SO_F)„. Contrasting magnetic D 6 z

behaviour is displayed by the fluoro-analogues of CsAg(S03F)3,

M^AgF^ where M1 = K, Rb, Cs which are reported to be antiferromagnetic, possibly by intermolecular superexchange via linear Ag2+—F —Ag2+ linkages. 183

TABLE 29

MAGNETIC SUSCEPTIBILITIES AND MAGNETIC MOMENTS FOR CsAg(S03F)3

T cor t XM yeff

(K) (106 cm3 mol"1) (y )

299 1564 1.91

274 1703 1.92

250 ' 1884 1.92

224 2095 1.92

199 2349 1.91

176 2673 1.92

152 3099 1.92

128 3714 1.92

108 4570 1.95

77 6146 1.90

aMagnetic moments are calculated using the Curie-Weiss law:

r 1/2 ueff = 2.828[xM° (T - 0)] 184

The electron spin resonance spectrum obtained on solid powdered CsAgtSO^F)^ showed an isotropic spectrum, most likely due to misaligned tetragonal axes rather than regular 0^ or

T^ symmetry, as in the case of AgCSO-^F^. Spectra recorded at

295 and 80 K give gQ-tensors of 2.182 and 2.184 respectively.

The magnetic moment of 1.89 u , calculated by equation (3.9), is in good agreement with the moments obtained from bulk magne• tic susceptibility measurements.

The electronic absorption spectrum of a neat thin powder film of CsAgCSO^F)^ was obtained along with its diffuse reflec• tance spectrum. Combining the two spectra results in an intense

UV band at 245 nm (40,800 cm ^); a strong broad band at 460 nm

(21,700 cm "*") , a shoulder at 560 nm (17,900 cm 1) and a weaker band at 720 nm (13,900 cm 1). Assignment of the three weak bands as due to d—d transitions in a tetragonally distorted

(elongated) octahedral environment with symmetry, consis• tent with the magnetic data and in analogy to the assignment proposed for AgtSO^F^ in Section III.B.3(b), would suggest an approximate 10 Dq value of 17,900 cm 1. The assignment:

2 2 1 2 2 1 B, -*- An at 13, 900 cm" ; B, —~ B„ at 17,900 cm" ; and lg lg lg 2g ' 2 2 -i Blg~~ Eg at 21,7 00 cm x with the expected splitting due to spin-orbit coupling unresolved; is comparable to those for

Ag^O^F^/ which has an approximate Ligand Field Splitting of

16,600 cm"1. In summary, the results of magnetic and spectroscopi measurements suggest CsAgCSO^F)^ to be a true divalent 2 + silver(II) compound with Ag in most likely a tetragonal elongated octahedral environment. IV. SILVER(II) TRIFLUOROMETHYLSULFATE

A. INTRODUCTION

Trifluoromethylsulfuric acid is comparable in acid

22 strength to fluorosulfuric acid, or slightly weaker as * * • 4- 210,211 ^ . . . . , . found in some studies . Originally prepared in

212

1954 by the oxidation of bis(trifluoromethylthio) mercury

and purified by treatment of barium bis(trifluoromethylsulfate) with fuming sulfuric acid, HSO^CF^ has attracted a lot of 22

interest, as evidenced in a recent detailed review article

on HSO^CF^ and its derivatives. HSO^CF^ is now manufactured

commercially by the electrofluorination of CH3S02C1 and

distributed as "Fluorochemical acid" by the Minnesota Mining

and Manufacturing Company.

Much like the fluorosulfate group, triflurormethylsulfate

can act as a monodentate, bidentate or even tridentate ligand

through oxygen. Discrete SO^CF^ anions exist in alkali metal

salts while in the titanium chloro salts the SO^CF^ group

actligans da s ian bidentatTiCl3S03eCF ligan3 d in TiCl2(SO^CF^)^ or a tridentate 79 22

In view of the interest in S03CF3 derivatives , it is 187

surprising that only a few transition metal trifluoromethyl- I 213 sulfates are known. These include Cu (SO^CF^) ,

11 213 105 11 105 Cu (S03CF3)2 ' ancj Co (S03CF3)2 with the metals 214

in rather common oxidation states; while Mo2(S03CF3)4 4 +

contains the quadruply bonded Mo2 ion. The scarcity of

examples is primarily due to the lack of suitable synthetic

routes. The very limited thermal stability of oxidizing agents, such as bis (trifluoromethylsulfuryl) peroxide,

215

(CF3S03)2 and chlorine(I) trifluoromethylsulfate,

216

CF3S03C1 restrict the available preparative methods to

solvolysis reactions in HS03CF3 with metal-halides, -carboxy-

lates, -oxides, and -carbonates as solutes. In some instances,

the method of ligand replacement by silver salt method (Section ' 217 I.C. 3) is also applicable, as AgS03CF3 is available commercially.

The purpose of synthesizing silver(II) trifluoromethyl-

sulfate was to investigate the possible stabilization of

silver(II) by yet another highly electronegative anion compar•

able to the fluorosulfate and more generally, to explore a

novel synthetic route to transition metal trifluoromethyl•

sulf ates . B. SYNTHETIC REACTIONS

1. Synthesis of Silver(II) Trifluoromethylsulfate

(a) Reaction of Ag(S03F)2 and HS03CF3

An excess of trifluoromethylsulfuric acid (^ 5 ml)

was distilled at reduced pressure (Section II.A.5) directly

onto a one piece pyrex reactor containing 1.012 g (3.31 mmol)

of Ag(S03F)2> The resulting suspension was magnetically

stirred at room temperature for two days. A slight pressure

increase was detected inside the reactor containing the

dark brown suspension, which was subsequently vacuum filtered,

and washed with five 0.5 ml portions of distilled HSO^CF^.

The remaining traces of volatile materials were removed in a

dynamic vacuum. 0.74 0 mg of dark brown, extremely hygroscopic

solid was obtained. The elemental analysis identified the

composition as Ag(S03CF3)2. Calculated: %Ag, 26.57; %S, 15.79

%F, 28.08; %C, 5.92. Found: %Ag, 26.81; %S, 15.92; %F, 28.04;

%C, 5.82.

The rather low yield, % 55 % based on moles of Ag is most

likely due to the method of isolation by filtration and washing. If the volatiles were removed simply by pumping under

vacuum, the resulting solid product was highly impure with a

mixed dark brown and white appearance, probably containing 189

Ag^O^CF^ as the impurity,

(b) Other Synthetic Attempts

(i) Reaction of Silver (II) Fluoride and HSO^F^

In a typical reaction, excess purified HSO^CF^ was added onto a sample of AgF., in a Kel-F reaction tube (Section II.A.3) inside the drybox. The resulting mixture was stirred at room temperature and the HF liberated was periodically removed by pumping on the mixture. Reaction times up to a week or even heating to 50 °C did not result in a homogeneous product. The volatile materials were removed as in the previous reaction.

Incomplete reaction is indicated by the carbon analysis of

4.31 % versus 5.92 % for Ag(S03CF3)2.

(ii) Reaction of AgF2 and (CF3S02)20

Excess trifluoromethylsulfonic anhydride was vacuum

distilled onto a sample of AgF2 in a Kel-F tube reactor. The expected reaction proceeded at room temperature,

2 (CF3S02)20 + AgF2 > Ag(S03CF3)2 + 2 FS02CF3

(4.1)

The gas phase IR spectrum of the volatile material showed the

217 218 presence of CF3S02F ' in addition to the starting material 190

(CFjSC^^0, Removal of all volatiles resulted in a slightly non-homogeneous dark-brown-black solid. The carbon analysis

indicated slighly impure Ag(S03CF3)2 with 6.20 % C versus

5.92 % calculated.

(iii) Oxidation of AgS03CF3 by S2OgF2 in HS03CF3

As AgS03CF3 was available commercially, oxidation with

S2°6F2 HS03CF3 was attemPted • Such a method has been 219

previously employed to synthesize I(S03CF3)3

Typically, a sample of AgS03CF3 dissolve in HS03CF3 in a one-part thick wall reactor was reacted with a small excess

of S20gF2 at ^ ice-water temperature. The reaction, once initiated, was highly exothermic. Allowing the reaction mixture to warm to room temperature without controlled cooling had resulted in an explosion, because the oxidative cleavage of the

S—C bond in HS03CF3 by S20gF2 is a recognized complication.

Subsequent filtration of the dark brown suspension and washing

with distilled HS03CF3 resulted in pure Ag(S03CF3)2 as supported by its carbon analysis (found 6.04 %) , the IR spectrum and thermal decomposition.

In summary it seems that none of the alternative methods

for synthesizing Ag(S03CF3)2 is as convenient and safe as

the solvolysis reaction of Ag(S03F)2 in HS03CF3.. The solvoly- 191

sis of transition metal fluorosulfates in HSO^CF^ as a synthetic route to the corresponding trifluoromethylsulfates appears to have general applicability as demonstrated by the synthesis of gold(III) trifluoromethylsulfate from its corresponding

4. fluorosulfat*i i c e 220

The successful quantitative conversion of a fluorosulfate to a trifluoromethylsulfate appears somewhat surprising.

Although this conversion seems to follow the general pattern of well known solvolysis reactions in HSO^CF^ with the general formula:

MX + n HSO,CF- »• M(SO-,CF_.) + n HX (4.2) n 3 3 jin

213,214 with X = Cl, 02CR or S03F, it differs from the others in two important aspects. In order to achieve complete conversion and to avoid mixed products, HX must be removed from the acid mixture, either:

(i) as a volatile byproduct, such as HC1 in the solvolysis of metal chlorides, or

(ii) as a soluble protonated species, as in the solvolysis of carboxylates.

For case(i), HS03F and HS03CF3 have virtually identical o 97 boiling points at atmospheric pressure, 162.7 C for HS03F o 212 versus 163 C for HS03CF3 , and similar volatilities under a dynamic vacuum at room temperature. Whereas in case (ii), a equally strong or stronger protonic acid, HSO^F, appears to be generated by a weaker protonic acid, HSO^CF^ from its salts

M(S03F)n- Hence removal of HS03F by virtue of its volatility

+ or by forming a protonated ion, such as H2S03F , is rather unlikely.

The possibility of removing the HSC^F by a chemical reaction is supported by a recent communication by Olah and 221

Ohyama , who refluxed a mixture of HSO^CF^ and HS03F at a

mole ratio of 2:1. CF3OS02CF3 and CF3S02F were obtained in 19 and 5.5 % yield respectively. No other reaction product was reported and the overall reaction is not completely understood

The reaction between HS03F and HS03CF3 is currently being 222

studied in detail . As HS03CF3 was used in large excess in

the solvolysis reactions compared to the amount of HS03F liberated, further reaction is very likely. The gas phase infrared spectrum of the volatile materials from the synthesis

of Cs[Au(S03CF3)4] from the reaction of Cs[Au(S03F)4] with

HS03CF3 shows band attributable to CF3OS02CF3, CF3S02F, F2CO, 220

S02 and SiF4 . Hence it may be concluded that Ag(S03F)2, o

the HS03F formed in the solvolysis reaction according to

equation (4.1), reacts with the excess HS03CF3 to form volatil

products, which do not interact with Ag(S03CF3)2. 193

2. Conversion of Ag(S03CF3)2 into [Ag(bipy)2](S03CF3)2

Chemical characterization of AgfSO^CF^)., is accomplished

by conversion into the previously published [Ag(bipy)2]- 113

(SO-jCF^),, . A similar synthetic procedure as used before

for the preparation of [Ag(bipy)2](SO^F) (Section III.C.2) was employed. After vacuum filtration, the resulting dark

reddish brown powder was identified as [Ag(bipy)2] (SO^CF^),, by the elemental analysis, infrared spectrum and melting point.

Found: %C, 36.17; %H, 2.31; %N, 7.47. Calculated: %C, 36.78;

%H, 2.25; %N, 7.80.

C. EXPERIMENTAL RESULT AND DISCUSSION

1. Infrared Spectrum

A rather complicated situation, due to the extensive mixing of CF^ and SO^ fundamentals is encountered for the

SO^CF^ ion. This is illustrated by two independent vibration• al studies of the SO^CF^ anion, including normal coordinate

223 224 analyses ' , which differ substantially on the vibrational 223 assignment. The assignment of Burger et al. will be used

in this discussion and comparison to spectra of previously 194

reported trifluoromethylsulfates will be used.

The infrared frequencies of AgtSO^CF^^ obtained from

neat solid film between BaF2 and KRS-5 plates are listed in

Table 30. A Raman spectrum could not be obtained.

When the band positions of Ag(S03CF3)2 are compared to 225

those of a bidentate SC>3CF3 group, as in (CH3) 3SnS03CF3 , reasonable agreement is obtained while the relative band intensities also match. Some shifts of the band positions

are to be expected when the S03CF3 groups are in an organotin environment. A large discrepancy in the relative band intensities is observed when comparing the spectrum with that 79

of the tridentate CF3S03 group in Cl3TiS03CF3 , even though the band positions matches well. Hence it appears that the

actual coordination mode of the S03CF3 groups in Ag(S03CF3)2 must be somewhere in between the purely bidentate and tridentate modes. Slight interaction between the third oxygen and silver(II) is possible.

2. Magnetic Measurements

The initial bulk magnetic susceptibility measurements

on Ag(S03CF3)2 gave somewhat surprising result. The sample was found to be antiferromagnetic with a rather low magnetic moment of 1.64 y at room temperature. The result was repro- 195

TABLE 30

INFRARED FREQUENCIES OF Ag (SO-,CF _ ) _ AND RELATED COMPOUNDS (cm )

3 Ag(S03CF3)2 (CH3) -^SnSO^F.^ Cl3Ti (S03CF3) Assignment

1280 vs 1319 vs, b 1272 vs vso3 (A")

1225 s, sh 1226 vs 1225 s, b VCF3 (Ai)

1 1202 vs 1179 s, br vS03 (A )

1125 vs 1145 s 1120 m VCF3 (E)

1026 1050 1035 s, b s m, sh 1 vS03 (A ) 1031 vs

775 w 771 m 780 mw vS-C (A1)

632 s 633 ms 630 s 6S03 bend (A")

595 m 595 n», sh 6S03 bend(A')

585 m 577 ms 6CF3 (E)

525 s 530 ms 525 ms 6S03 bend(A') 517 ms 490 m 452 vs TiO str. 428 vs TiCl region

365 mw 356 ms 391 m, sh 6C-S-0 (A")

335 mw 347 s 6C-S-0 (A 1)

325 vw , sh 317 m 6CF3 (A1) 196

ducible, with freshly prepared samples giving identical results.

Hence impurities or partial decomposition as possible causes must be ruled out. In addition, elemental analysis subsequent

to magnetic measurements confirmed the purity of the sample.

(Calculated: %Ag, 26.57; %C, 5.92. Found: %Ag, 26.56; %C, 5.81)

The result of the magnetic measurements is shown in

r Table 31. A plot of XM° versus T results in a Neel tempera•

ture of 138 K.

The dipolar coupling approach used in Section III.D.3. is

again attempted to determine the extent of the exchange inter•

action. Best agreement between theoretical plot and experimen•

tal data is obtained for the n = 2 plot. The agreement between

the average and individual J values is good, except at the lower

limit of the temperature range, as shown in Table 32. The

average J value is found to be 84.9 cm 1.

Knowing the Neel temperature, T , from the experimental

data, the value of the exchange integral can be estimated.

The magnetic susceptibility of a two center exchange system

225 with S = 1/2 is found to be :

-1 [ 1 + i exp( - 2 J/kT ) ] (4.3) X M 3 kT

Let g = 2 as a first approximation and = 1/8 3 k TABLE 31

MAGNETIC SUSCEPTIBILITY DATA OF Ag(SO,CF.J

T % . P6ff (K) (106 cm3mol-1) (u )

304 1099 ± 17 1.64 ± .02

278 1148 ± 6 1.60 ± .02

254 1203 ± 7 1.56 ± .02

226 1268 ± 11 1.52 ± .02

204 1324 ± 18 1.47 ± .02

179 1381 ± 21 1.41 ± .03

154 1421 ± 27 1.32 ± .02

128 1425 ± 31 1.21 ± .02

108 1390 ±40 1.10 ± .03

76 1421 ± 31 0.93 ± .02 198

_1 XM = [ 1 • + i exp ( - 2 J/kT ) ] (4.4) 2 T

Setting = 0 in equation (4.4) at T would lead to

226 J/kTN - - 4/5 . With T^ = 138 K for Ag(S03CF3)2, the exchange integral J is estimated to be approximately 77 cm-1.

Comparing to the value of 84.9 cm 1 obtained from the plot, a reasonable agreement within 10% is obtained.

Based on the discrepancy in the magnetic behaviour between

Ag(S03F)2 and Ag(SC>3CF3)2, it must be concluded that metal- metal direct interaction is feasible in the trifluoromethy1sul• fate, but not in the fluorosulfate, possibly due to a difference in coordination mode. It may be speculated that the bulkier

CF3S03 groups impede the involvement of the third oxygen in coordination, thus allowing direct Ag—Ag interaction in the solid state.

3. ESR Spectrum

The ESR spectrum of Ag(S03CF3)2 was recorded at 295 and

80 K. Broad isotropic signals were obtained at both tempera• tures. The g-tensors obtained were 2.175 and 2.162 at 295 and

8 0 K respectively, comparable to other Ag(II) containing compounds. Although magnetic moments could be calculated from

the gQ values using equation (3.9), results so obtained may not 199

TABLE 32

EXPERIMENTAL J VALUES OF Ag(S03F)2

n = 2

expt kT kT/J yeff

-1, 1 (yB) (cm (cm )

1.64 ± .02 211. 3 •2.52 ± .18 83.8 ± 6.4

1.60 ± .02 193.6 •2.17 ± .17 89.2 ± 7.6

1.56 ± .02 176. 2 •1.93 ± .10 91.3 ± 5.0

1.52 ± .02 157.4 •1.76 ± .08 89.4 ± 4.3

1.47 ± .02 141.4 •1.59 ± .08 88.9 ± 4.7

1.41 ± .03 124 . 4 •1.42 ± .08 87.6 ± 5.2

1.32 ± .02 107.0 •1.24 ± .03 86.3 ± 2.1

1.21 ± .02 89.3 •1.06 ± .03 84.3 ± 2.4

1.10 ± .03 75.1 •0.94 ± .04 79.9 ± 3.6

0.93 ± .02 53.2 •0.78 ± .02 68.2 ± 1.8

-1 J = 84.9 cm 200

be very meaningful in view of the antiferromagnetic behaviour.

As discussed earlier, the antiferromagnetic Ag^SO^F)^ and

K2Ag(S03F)4 both gave broad single lines, while the earlier reported antiferromagnetic silver(II) bis(nicotinate) also gives a broad resonance centered at g = .2.08 in the ESR 144 spectrum

4. Thermal Decomposition

Under one atmosphere of dry nitrogen, Ag(SC<3CF3)2 is thermally stable up to 14 0 °C, at which it decomposed into a white solid with gas evolution. To study the decomposition

in detail, a sample of Ag(S03CF3)2 (0.180 g, 4.42 mmol) in a one-piece pyrex reactor was heated at 150 °C for one hour.

The volatiles evolved was condensed into a storage container at - 196 K. No incondensable gas was detected. The infrared spectrum of the white solid residue was found to be identical

to the spectrum of the commercial AgS03CF3 obtained from

Ventron Alfa Corp. except for a single weak band at 8 00 cm ^.

The weight of the residue was 0.116 g as compared to 0.114 g

1 expected for Ag (S03CF3).

A gas phase infrared spectrum was obtained on the trapped volatile materials. The absorption maxima of the complicated

spectrum are listed in Table 33. Most of the bands can be 201

TABLE 3 3

IR FREQUENCIES OF VOLATILES

FROM

THERMAL DECOMPOSITION OF Ag(S03CF3)2

(cm 1)

227 202 215 228 70 Ag(S03CF3) 2 COF. S02F2 CF3S03CF3 SO, SiF/ (decomposed volatiles)

1930 m, b 1928 1500 w, sh 1502 1465 vs 1461 s 1355 w, sh 1362 1290 1269 1258 s vs 1249 b 1230 1230 s 1150 s, sh 1151 1135 vs 1134 s 1030 w 1010 960 vs 965 954 s 885 848 w 848 790 s 786 m 776 ms 774 766 m 766 m 715 w 626 s 626 605 s 585 m 584 544 542 w 545 512 w 518 390 w, sh 386 390 202

satisfactorily assigned when compared to the reference spectrum

221 202 222 228 17 of COF2 , S02F2 , CF3S03CF3 , S02 and SiF4 °, as shown in Table 33. It is plausible that the thermal

decomposition of Ag(S03CF3)2 resembles that of Ag(S03F)2 ini• tially, with subsequent thermal decomposition and rearrangement

215 of the S03CF3 radicals .

D. CONCLUSION

The conversion of Ag(S03F)2 into Ag(S03CF3)2 opens up another synthetic route to transition metal trifluoromethyl•

sulf ates, via solvolysis reactions in HS03CF3. This is

further supported by the analogous conversion of Au(S03F)3 and

Cs[Au(S03F)^] into their corresponding trifluoromethylsulfates

As the existing transition metal fluorosulfates and the avail• able synthetic routes to fluorosulfates outnumber those of

the trifluoromethylsulfates, the scope of the possible S03CF3 containing compounds is widened and synthesis of compounds with the metal in an unusual oxidation state, such as

Ag(S03CF3)2, is now feasible.

The stabilization of silver(II) by fluorosulfate and also trifluoromethylsulfate suggests the possibility of synthesizing other binary silver(II) compounds of highly electronegative anions, provided suitable synthetic routes are available. 203

V. FLUOROSULFATES OF RUTHENIUM

A. INTRODUCTION

Both ruthenium and osmium exhibit in their compounds formal oxidation states up to +8; the latter being the highest for all d-block transition metals, is represented by RuO^ and OsO^. The non-existance of the corresponding octafluorides 229 may be due to steric reasons . The highest binary fluorides 14 of these elements are RuF, and the rather unstable OsF_ 6 7

From the experience gained from the silver system, it appear• ed rather unlikely that binary fluorosulfates or fluorosulfato complexes of osmium and ruthenium would exist with the metal in such high oxidation states. It appeared however interest• ing to investigate the possible fluorosulfates of these two elements for the following reasons: (i) to test the synthetic routes developed in the study on divalent silver compounds;

(ii) to conduct a study of hopefully magnetically dilute compounds with the metal in the intermediate oxidations states

(perhaps +5 to +3); and (iii) to investigate the spectroscopic properties of these compounds.

The earlier work on the chemistry of ruthenium and the other rarer Group VIII metals, Os, Ir and Rh, has been describ- 204

230 ed by Griffith . Of all the oxidation states displayed

by ruthenium, the trivalent state is most commonly found.

As expected, the chemistry of ruthenium resembles that of

osmium far more than iron. The greatest similarities are

found in the higher oxidation states (VIII to V), which are

dominated by oxides and fluorides; and also in the lower oxida•

tion states (II to 0), by the numerous complexes with ir-acid

ligands. The chemistry of ruthenium and other 4d, 5d

Group VIII metals in the lower oxida'tion states is a very

active area of research, mainly in the field of homogeneous

catalysis.

The halide chemistry of ruthenium will be briefly

reviewed. The existing binary fluorides of ruthenium have

been shown in Figure 2 to be +6, +5, +4 and +3. Ruthenium

hexafluoride, RuFg, is made by direct fluorination of the metal 2 31 to give dark brown crystals . The infrared spectrum of the 2 32 vapour confirms the suggested octahedral geometry . The +6 oxidation state is also found in the oxyfluoride RuOF^, a para• magnetic (u .p- = 2.91 u_ at room temperature) pale green err a 233

crystalline material ; and in Cs2[Ru02Cl4], made by the 234

reaction of HC1 and RuO^ in the presence of CsCl

Ruthenium pentafluoride is a cis-fluorine-bridged tetramer.

This dark green solid is best prepared by the direct reaction of the metal and fluorine at temperature between 280 and 233 900 °C . The vigorous reaction of BrF^ and the metal 205

yields a RuF,.—BrF^ complex along with RuOF4. RuF,., often with

traces of RuOF4,is obtained on decomposition of this initial 235 complex . RuFj. has a room temperature magnetic moment of 233

3.60 y , very similar to 3.48 y for K[RuF ] , suggesting little metal-metal interactions. RuF5' i-s a good fluoride ion acceptor. Examples may be found in the chemistry of noble

gases where complexes of the type XeF-^'MF,., 2XeF2 *MF,- and

XeF2*2MF5 can be prepared by the direct reaction of the con- ... . j., ... 237,238 stituent fluorides K [RuFg], can be made by the reaction of BrF^ with a mixture 235 of potassium bromide and ruthenium in a 1:1 mole ratio However, the interaction between Ru and BrF^ is reported to be 12

explosively violent , and should be moderated by using Br2 as a diluent. Other hexafluororuthenates(V) have other alkali metals, Ag(I) or Tl(I) as counter cation. The electronic spec• trum of the solids between 10,000 and 40,000 cm 1 have been 239 240 241 measured and assigned to d—d transitions ' '

Ruthenium tetrafluoride is obtained from the reaction of

RuF,. with I2 and IF,. . These yellow crystals has a room tempera- 242 ture magnetic moment of 3.04 yD . RuCl. has been identified in the vapour phase during the chlorination of RuCl^ at 750 °C, 243 244 but its existence in the solid state seems doubtful ' 2-

Hexahaloruthenates(IV) of the type [RuXg] , where X = F, Cl, or

Br have been synthesized with alkali metals as counter cations. 206

mi -» , , , 239,245 ^ . ,241 The electronic spectra have been measured and assigned

The magnetic susceptibilities of the fluoro-and chloro- 246 247

complexes have been studied in detail ' . Magnetic moments

close to spin-only values at room temperature have been found

for these magnetically dilute compounds.

Ruthenium trifluoride can be made by the reduction of the

pentafluoride with excess iodine at 150 °C 2^8. Neutron

diffraction powder patterns taken at 4.2 K and also 298 K 24 9 suggest that this brown compound is not magnetically ordered Hexafluoro-complexes of ruthenium(III) of the type M^RuF^ have 241 been synthesized and the electronic spectra studied . In particular, the dark grey K^RuF^ was prepared from the fusion 12

of RuCl^ with KHF2 in a stream of nitrogen . Lower fluorides

of ruthenium do not exist.

Ruthenium trichloride is commonly used as a starting material

for many higher and lower valent ruthenium compounds. Anhydrous

RuCl^ has two modifications. Interaction of the metal with a

mixture of CC>2 and Cl,, at 330 °C gives the dark brown 3-form, which can be irreversibly converted to the black leaflets of

250 251 the a-form at ^ 550 °C under Cl2 ' > The structure of 3-RuCl^ has linear chains of Ru atoms in —RuCl^—RuCl^—RuCl^— 3

units with a P3cl (C^v) space group whereas a-RuCl^ has a P3112 3

(D^) space group with a distorted octahedral environment around

ruthenium. In summary, in binary fluorides of ruthenium, the higher oxidation states are preferred. Octahedral coordination around

I the metal, as in RuF,, M (RuF,), (RuFr)., M*(RuF,), is often D D O 4 O observed. Intermediate and lower oxidation states in binary compounds are found in chlorides, bromides, and iodides. No fluorosulfates of ruthenium have been reported in the literature

B. RUTHENIUM(III) FLUOROSULFATE

1. Preparation and Elemental Analysis

The metal powder is a suitable starting material, as it is

available in high purity. The reaction with HS03F/S20gF2 was carried out in a similar manner to the silver metal reactions.

A one-part pyrex reactor containing 0.253 g (2.50 mmol) of ruthenium powder was flame dried in vacuo to remove residual moisture. About 5 ml of HSO^F were distilled directly into the reactor and about 2.5 ml of S~0,F„ were then added in vacuo. 2 b 2

The mixture was allowed to warm to room temperature and then magnetically stirred at 60 °C for one day. A dark red-brown solution resulted. Removal of all volatile materials was accomplished by warming the reactor and its content to 6 0 °C in vacuo. An almost black, dark red-brown solid remained. [1.05 0

as compared to 0.995 g expected for Ru (S0oF) ] . Elemental 2 08

Analysis confirms the composition, calculated: %Ru, 25.38;

%S, 24.15; %F, 14.31. Found: %Ru, 25.15; %S, 24.02;

%F, 14.46. Ru(S03F)3 is hygroscopic and thermally stable up

to 14 0 °C, above which it decomposes to a red brown liquid.

2. Experimental Result and Discussion

(a) Infrared Spectra

Infrared spectra of Ru(S03F)3 were recorded on thin solid

films between BaF2, KRS-5, AgBr and also AgCl plates. The

spectra recorded with different window materials are identical.

The resolution is generally poor in the S—O stretching region,^

because of the rather broad band contours. A Raman spectrum

could not be obtained due to the dark sample color. The IR

absorption frequencies are listed in Table 34.

The tentative assignment in Table 34 suggests the presence of vibrations due to bi- and mono-dentate fluorosulfate strong resemblance is observed when the above spectrum is compared to 70

the published IR spectrum of Sn(S03F)4 , as shown in Table 34, The presence of mono- and bi-dentate SO^F groups in the poly- 119 meric hexa-coordinated Sn(S03F)4 is supported by its Sn .. , 70 Mossbauer spectrum 5 6 2

On the other hand, the high spin d Fe(S03F)3 shows IR absorptions (also listed in Table 34) typical of pure bidentate

fluorosulfates, consistent with regular octahedral coordination around Fe^+. 209

TABLE 34

INFRARED FREQUENCIES OF Ru(SO,F), AND RELATED COMPOUNDS (cm )

/u Ru(S03F)3 Fe(S03F)3 Sn(S03F)4 Assignment

1430 s, sh 1438 s v SO„ (mono, A") asy 2 1411 s

1390 vs, b 1360 m 1395 w, sh v so3 (bi, A")

0 1220 s 1232 S V S0 (mono, A') sym 2

1110 sh 1137 s 1130 s v S03 (bi, A')

1025 vs, b 1090 s 1085 s v S03 (bi, A') 980 w, sh

920 w 920 m v S-O (mono, A 1)

860 s 850 m 850 m v S-F (mono, bi) 820 sh 832 s

645 mw, b 630 m 657 m, sh 605 w 628 m

585 m 579 w 585 m 550 ms 551 m 550 s

475 vw, sh 460 w 450 w 442 w 438 m

395 vw 419 w

328 m

300 w, b 318 m 294 w 210

(b) Magnetic Susceptibilities

Due to the larger ligand-field splittings of the second row transition elements, spin pairing is expected. The elec• tronic configuration for ruthenium(III), 4d^, is expected to

result in a T2g ground term for an octahedral environment.

In fact, all magnetically dilute ruthenium(III) compounds studied are of the low-spin type. 2 Magnetic moments for ions with T~ terms have been cal- 2g 2 52 culated and plotted against kT/A where A is the spin-orbit coupling constant of the ground term. A may be obtained from

the free ion spin-orbit coupling constant, Aq, to give a value -1 253 of - 118 0 cm . The magnetic moment expected should be between 2.0 and 1.7 y_ in the temperature range of 300 and 80 K respectively. The observed magnetic moments for most octahedral 90,254 complexes of ruthenium(III) fall into this range with the 2 55 following exceptions: (i) the dimeric [RuCl3(Et2S)2]2 , where the low room temperature magnetic moment of 0.95 y_ per B Ru has been explained by spin-coupling between the ruthenium 25 6

ions in the black solid; (ii) in K2RuCl5 , where the variation

of yeff from 1.64 to 1.14 yg between 300 and 80 K has been suggested to be due to magnetic exchange between the ruthenium atoms in the polymeric compound; and (iii) the brown 3-form of

RuCl^, which is antiferromagnetic below ^ 600 K and has much lower

„ , . , ., ., ... j. 257,258 cor , m magnetic susceptibilities than the a-form Y at T M N 211

is estimated to be 170 cm mol (^0.9 y ) and falls to zero at

^ 150 K. The black a-form of RuCl^ obeys Curie-Weiss Law above

^ 50 K with 0 = 40 K, but is antiferromagnetic at lower temperatures with a Neel temperature, of 13 K. The magnetic moment is 2.17 y at 293 K for the a-form.

B The corresponding FeCSO^F)^ has a magnetic moment of 6 2

5.33 yB at 298 K , and is reported to show some temperature and magnetic field dependence. Nevertheless, the presence of 5 the expected hxgh spin d ions in FeCSO-^FJ^ is confirmed.

The magnetic susceptibility of RuCSO^F)^ was measured between ^ 300 and 77 K. The resulting susceptibility increases with temperature with a room temperature magnetic moment of

^ 0.97 yB. The measurements were repeated with a freshly prepared batch of sample with almost identical results. The average susceptibilities data is listed in Table 35.

Unlike the 6-form of RuCl^, the rather low magnetic moment of RuCSO^F)^ is fairly independent of temperature. In addition, the extra Ligand Field Stabilization Energy with an octahedral coordination would make a T^ symmetry rather unlikely. The pre• sence of direct interaction between the ruthenium atoms through metal-metal bonding is possible. Such an interpretation is compa• tible with the existance of both mono- and bi-dentate fluorosulfate groups discussed earlier. The almost black color of RulSO^F)^ is often observed in compounds with exchange interactions, such as 212

TABLE 35

MAGNETIC SUSCEPTIBILITY DATA OF Ru(S03F)3

cor ^M u T u yeff 1 yeff

6 3 _1 (10 cm mol ) (yB) (K) (uB)

391 0.966 298 1.03

410 0.946 273 1.02

437 0.934 250 1.01

477 0.922 223 1.01

'524 0.909 197 1.00

576 0.897 175 1.00

656 0.890 151 1.01

745 0.871 127 1.01

884 0.876 108 1.04

1120 0.830 77 1.04

C = 0.1291 ± 0.0036

0 = - 43.4 ± 6.7 the anionic silver(II) complexes M^AgCSO^F)^ discussed earlier 255

and also the dimeric [RuCl^(Et2S) ]2 . Application of the graphical dipolar coupling method to determine the number of exchanging centers, as attempted for the silver compounds is unfeasible here, as the variation in magnetic moment is only large enough to make use of two or three of the theoretical 193 data points provided

The powder diffuse reflectance spectrum of RuCSO^F)^, recorded between 350 and 730 nm, shows only a single broad band centering at ^ 450 nm. Since RuCSO^F)^ is soluble in HSO^F, as evidenced from the synthesis of the compound, the spectrum of a 0.294 mol/1 solution in HSO^F was measured at 1/10 and

1/100 dilutions. The resulting spectrum is illustrated in

Figure 21. Based on the molar extinction coefficients, the three highly intense bands at lower wavelengths: 447 nm

(22,400 cm-1, e = 3880), 310 nm (32,300 cm"1, e = 3570);, and

220 nm (45,500 cm-1, e = 8440), are most likely of the ligand metal charge-transfer type, whereas the weak band at 735 nm

(13, 600 cm = 95) appears to be a d—d transition judging by the low molar extinction coefficient. 258 24 5 J0rgensen ' has investigated the solution spectrum 3- of [RuCl,] in 12 M HC1. The spectral information is listed b FIGURE 21

—| , 1 1 1 200 300 400 500 600 nm 215

in Table 36 along with those of RuCSO^F)^ and other related complexes. When comparing the general band shape of the 3-

[RuCl&] spectrum to the solution spectrum of Ru(S03F)3 in

HS03F, some resemblance is apparent. The major difference is the shifting of the weakest band to lower energy for

Ru(S03F)3. J0rgensen has assigned all the higher energy bands above ^ 20,000 cm 1 to charge-transfer transitions except the lowest energy absorption at 19,000 cm 1 (e = 40), attributed to a spin-allowed t^ —» t^ e transition 2^8. ^ 2g - 2g g Oh the other hand, the solid state electronic spectra 5 of some 4d hexafluorometallate complexes of Ru(III) .and 239 241

Rh(IV) have been studied in two instances ' . As shown in Table 36, bands observed below ^ 4 0,000 cm 1 have all been assigned to spin-forbidden and spin-allowed transitions within the metal d-orbitals. Nevertheless, the overall feature of 3_ the spectra for [RuCl^] and Ru(S03F)3 in HSO^F appear similar to that of the diffuse reflectance spectrum of K^tRuFg] 241 reported more recently by Allen, et al. . They suggest that the charge transfer bands of the fluoro-complex should appear at substantially higher energies than the other halo- 259 complexes based on electronegativity considerations

In view of the high extinction coefficients of the higher energy bands in the solution spectrum of RutSO^F)^ in HSO^F, TABLE 36

ELECTRONIC SPECTRA OF Ru(SO-F) AND RELATED COMPLEXES (x 103 cm"1)

1_ TCO T OQ 941 Ru(SO-,F)_ RuCl, Cs-RhF, K-.RuF, 3 3 b 2 6 3 6 in HSO^F in 12M HC1 Assignment powder Assignment powder Assignment

12.2 2T - 10.0 2g - s 2g \ ig

4 1 2 13.6 19 . 0 t e 16.1 15.4 T - 2g 2g g 2g 2g 2g S (95)a (40) 2g 2 2 2 22.4 25.6, sh 7T — 19-21 20.0 2g - \ A T T 2g 2g 2g, 2g' lg' 2g (3,880) (600) \ 2E g

2 2 2 28 .7 TT — 26.0 . 2T - 26.5, sh ZE /T. /T_ fc2g 2g ig 2g g ig 2g (2,200)

32 .4 TT — 34.0, br T„ - 32.3 fc 2g 2g ig (3,570) (1,700) to 45.5 43.6 IT — >50 .0 TT — fc2g 2g (8,440) (16,000) l -1 Dq = 2100 cm" Dq = 2200 cm

a. e(M cm ), molar extinction coefficients in brackets. to attribute them to spin-allowed d—d transitions would be inappropriate. Hence it appears that,except for the observed band at 13,600 cm 1, the other spin-allowed d—d transitions

RuCSO^F)^ are mixed-in or hidden by the much stronger charge transfer absorptions. Deduction of the ligand field splitting parameter, 10 Dq, is therefore not possible here. A further difficulty arises from the fact that the exact nature of

Ru(S03F)3 in HSO^F is uncertain, even though the spectral 3- similarity to complexes of the type RuX^. suggest a similar species in HSO^F solution. However, attempts to isolate 3- complexes containing the RuCSO^F)^ anion out of solution have not been sucessful. This will be further discussed later on.

(d) ESR Spectra

The electron spin resonance spectrum of Ru^O^F)^ powder was measured at 295 and 80 K while the same RuCSO^F)^ in HSO^F solution used in recording the electronic spectrum was measured at 80 K. The data obtained is listed in Table 37.

Extremely broad single lines (line width ^ 1000 gauss) that show some anisotropy were obtained from the solid sample, but parallel and perpendicular components of the g-tensor were not resolved. On the other hand, the HSO^F solution gave an aniso• tropic spectrum with three distinct principle g-values. 218

TABLE 37 ESR DATA OF Ru(S03F)3

Sample T(K)

Ru(S03F)3 powder 295 gQ = 1.804

Ru(S03F)3 powder 80 g = 1.777

Ru(S03F)3 in HS03F 80 gx = 1.515 = gj

g2 = 1.831

g3 = 1.956

(go = 1.777)

99 101 Hyperfine coupling due to Ru and Ru, both of nuclear spin

I = 5/2, was not detected. This is perhaps due to the rather low 92 natural abundance of 12.81 % and 16.98 % respectively for the two isotopes. Few studies on the ESR spectra of Ru(III) compounds have 92 3+ been published . The ESR spectrum of Ru in K3InClg*2 H^O gives 1.0, 1.22 and 3.24 as the three components of the

3+ g-tensor while in the case of Ru in Al(acac)3, g-values 261 262 of 1.28, 1.74 and 2.82 are obtained . One elaborate study

n c ^5 , c ^ 3+ 3+ T 4+ ^ , of many low-spin d complexes of Ru ,0s , Ir concentrates on complexes of low symmetry, the Ru3+ complex with the highest

+ symmetry (D4h) studied is [PEt3H] [RuCl4(PEt3)2]~ in frozen solution, which has 1.64, 2.51 as the gj | and g^ values.

The unresolved anisotropy observed in the powder spectrum of

Ru(S03F)3 at both temperatures appears to indicate a reduction of regular octahedral geometry around ruthenium, which is consistent with the presence of metal-metal interactions and two different types of coordination for the fluorosulfate groups. 219

(e) Discussion

To explain the low magnetic moment for Ru(S03F)3, a binuclear structure with multiple Ru-Ru bond similar to the groups of binuclear ruthenium carboxylates of the general formula

263,264 Ru(OCOR)4Cl where R - Me, Et, n-Pr could be considered. 263

The structure of Ru2(OCOC3H7)4C1 is illustrated in Figure 22

The proposed analogous [Ru(S03F)3]2 would have four bridging

bidentate SC»3F groups linking the Ru-Ru bonded dimeric units,

with two terminal monodentate S03F groups bonded along the

Ru-Ru bond axis. The paramagnetism (Peff between 2.76 and

3.37 uB) of Ru2(OCOR)4Cl is rationalized by a Molecular Orbital

diagram, suitable also for M2Xg species with D^ symmetry (e.g. 2-

Re2Clg ), with three unpaired electrons in each of the * non-bonding a, (a ), a- (a') and the antibonding b, (6 ) orbitals ^ lg n 2u n ^ lu two as shown in Figure 2 3 . For [Ru(S03F) J^* unpaired electrons are expected to be present in the two lower energy orbitals, a,l g and a„2 u . However,r the magneti? c moment calculated based on

[Ru(S03F)3]2 varies between 1.37^ at 2 98 K and 1.01 yQ at 77 K per dimeric unit, which is too low for two unpaired electrons. Such

a low Ve£f value is also inconsistent with the formulation as a mixed valency Ru11—RuIV system, where two unpaired electrons 4 are expected for the low spin d centers.

Alternatively, the low magnetic moment observed for

Ru(S03F)3 may be explained by (i) weak direct metal-metal interactions resulting in partial spin-pairing, since FIGURE 2 2 FIGURE 2 3

263 THE STRUCTURE OF Ru2 (O^C^) Cl 263 MOLECULAR ORBITAL DIAGRAM FOR Ru2(O^C^) Cl

(derived from M2Xg species with D4h symmetry)

blg(a )

e (TT*)

blu(6 )

a2u(aA} alg(an)

4f b2g(6)

e (TT) u 221

antiferromagnetic coupling through bridging SO^F groups would

imply the rather unlikely coupling through four bonds; (ii) the

simultaneous presence of magnetically dilute paramagnetic centers

of Ru(III) and diamagnetic binuclear Ru-Ru bonded pairs. The

experimental magnetic susceptibility was found to show a

Curie-Weiss relationship with a Weiss constant of - 43.4 K.

The magnetic moment calculated using equation (3.2) is C W

= independent of temperature, with Vejf" 1.02 ± 0.02 yB as

shown in Table 35. However, the use of equation (3.2) must

be viewed with caution since 0 is rather large here. Curie-Weiss Law behaviour is consistent with a non-octahedral environment 3+ 2

for Ru . The T2g ground term, expected for a regular octa•

hedral geometry will be replaced by a singly or doubly

degenerate term depending on the type of distortion.

Using equation (3.9), the experimental gQ values from the

ESR spectrum of RutSO^F)^ powder may be converted to a magnetic moment of 1.56 y at 295 K and 1.54 y at 80 K, which is not

unexpected for Ru3+. If case (ii), the presence of magnetically dilute Ru3+ is assumed, the observed ESR signal may be attribu• ted entirely to this species based on the rather insensitive nature of the signal to temperature changes, which is unlikely 194

for spin-paired triplet states . Such an assumption would allow an estimate of approximately 33 % of magnetically dilute

centers in Ru(S03F)3. However, a definite conclusion will require a crystal structure determination. 222

C. ANIONIC FLUOROSULFATO COMPLEXES OF RUTHENIUM(III)

1. Introduction

The solubility of the probably polymeric ruthenium(III)

fluorosulfate in HSO^F suggests the possible breakdown into

perhaps monomeric, solvated RuCSO^F)^. It follows that,

n anionic complexes with the general formula [Ru(SO^F)3+n]

should form. The use of chloryl fluorosulfate, C102S03F, as 42 172

a fluorosulfate donor has many precedents ' . The chief

advantages are : (i) the readily formed chloronium cation,

Cl02+,is detectable by its vibrational spectrum; (ii) when used

in excess, its solvating and ionizing ability does not require

the use of another solvent; (iii) its low volatility still

allows vacuum transfer of excess amounts. Hence it is generally

easy to detect the formation and the stoichiometry of anionic

complexes, as illustrated by the reaction of C102S03F and

Ag(S0.jF)2 discussed previously.

2. Synthetic Reactions and Elemental Analyses

(a) C102[Ru(S03F)4]

Excess dark red ClO^O^F (^ 5.6 g) was added to a one-piece

pyrex reactor containing 0.925 g (2.32 mmol) of RuCSO^F)^ inside

the drybox. After stirring the mixture at ^ 70 °C for one day,

complete dissolution of Ru (SO-,F)-. was observed. Excess ClO-SO^F 223

was then removed in vacuo with the reactor and content kept

at 70 °C. Due to the low volatility of C102S03F, the residual amount was pumped off with the reactor at 100 °C, leaving

1.334 g of black solid behind. The expected weight for

C102[Ru(SO^F)4] was 1.312 g. The elemental analysis provides confirmation. Calculated: %Ru, 17.90; %C1, 6.28; %F, 13.46.

Found: %Ru, 17.68; %Cl, 6.38; %F, 13.58.

(b) Cs[Ru(S03F)4]

A 1:1 stoichiometric mixture of Ru(SC>3F)3 (1.739 g,

4.366 mmol) and CsS03F (1.013 g) was loaded into a one-part

pyrex reactor. Excess HS03F ( ^ 10 ml ) was directly distilled onto the mixture. Subsequent stirring at ^ 60 °C for one day

resulted in complete dissolution. The HS03F was then removed in vacuo at ^ 6 0 °C. The homogeneous very dark brown-black solid obtained was identified by its elemental anlaysis to

have the composition Cs[Ru(S03F)4]. Calculated: %Ru, 16.04;

%Cs, 21.09; %F, 12.06. Found: %Ru, 15.90; %Cs, 20.97;

%F, 11.96.

3 Even though the existence of the Ru(S03F)g ion in the

solution of Ru(S03F)-? in HSO^F is suggested by the electronic 224

spectra as mentioned , the reaction between excess CIC^SO^F and RulSO^F)^ did not result in a hexakisfluorosulfatoruthenate(III) complex. Lower thermal stability of such a complex is a possible reason. An attempt was made to synthesize

such a complex using a 3:1 mole ratio of CsSC>3F and Ru(SC>3F)3.

A similar procedure as that employed to synthesize

Cs[Ru(S03F)4] was used. After removing the solvent (HS03F)

from the solution containing 0.898 g of Ru(S03F)3 and 1.569 g

of CsS03F, a non-homogeneous white and dark brown-black preci• pitate was obtained. The infrared spectrum of this product

shows a composite spectrum of CsS03F and Cs [Ru (SC>3F) ^ ] . It

appears that even with an excess of CsS03F, Cs [Ru (SC>3F) ^ ] is the only species formed.

3. Characterizations

(a) Infrared Spectra

Infrared spectra of neat powdered films of Cs [Ru (SC>3F) ^]

and C102[Ru(S03F)4] were measured between KRS-5, AgBr and

BaF2 window plates. The absorption bands are listed in Table 38.

The IR spectrum of Cs[Ru(S03F)4] is shown in Figure 24. No

Raman spectrum could be obtained on these dark colored samples.

The spectra of the two complexes are very similar except for

+ the vibrations due to C1C>2 cation in the spectrum of

ClO^[Ru(SO^F)A]. The band positions, in particular in the 225

TABLE 38

IR FREQUENCIES OF ANIONIC FLUOROSULFATO COMPLEXES OF RUTHENIUM(III)

(cm 1)

Cs[Ru(S03F)4] (C102)[Ru(S03F)4] Assignment

1360 s, b 1360 s, b V S02 (mono, AV) asy + 1295 m V C102 (v3) asy 1 1195 vs 1190 vs V S'02 (mono, A ) sym

1150 sh 1140 sh V S03 (bi, A')

+ ^1040 sh V cio2 (v^ sym 960 s, vb 940 s, vb V S—O (mono , A1 )

795 s, b 810 s, b V S-F

625 w 630 w

575 m 575 m 545 m 545 m

+ 515 w 6 C102 •(v2)

475 vw 475 vw 435 vw 435 vw

285 w, sh 260 w NJ NJ 227

SO^F stretching region, are characteristic for monodentate fluorosulfate groups in an anionic environment. A shoulder at ^ 1150 cm 1 is generally indicative of bidentate fluorosul• fate groups, however other characteristic bands at ^ 108 0 era"1 are not observed. The observed spectra of both complexes differ from the composite spectra of the starting materials,

in particular in the case of Cs[Ru(SO^F)4], where the absorptions

due to the SO^F 82,74 are a>osent> This observation argues against the presence of simple mixtures of the starting materials.

(b) Magnetic Susceptibility

The magnetic susceptibility of Cs[Ru(SO^F)4] was measured between 300 and 77 K. The resulting data is listed in Table 39.

Weak paramagnetism was observed with the magnetic moment decreasing with temperature from ^ 0.54 to ^ 0.31 y^. Like

RuCSO^F)^, the low magnetic moment observed here may again be due to metal-metal interaction. However, the susceptibility measured is much smaller and does not show a Curie-Weiss relationship (C = 0.080 ± 0.017). Further information is expected from the ESR measurements discussed below.

Measurements at room temperature gave barely detectable signals, hence no quantitative result could be obtained. 228

TABLE 39

MAGNETIC SUSCEPTIBILITY DATA OF Cs[Ru(SO..F) ]

cor m *M ^eff

6 1 (10 cm^ol" ) (yB) (K)

120.4 0.538 300

135.2 0.545 275

139.1 0.527 250

141.6 0.505 225

145.9 0.483 200

154.2 0.466 176

164.8 0.466 151

171.6 0.419 128

177.6 0.387 106

151.1 0.306 77 At 80 K, C102[Ru(S03F)4] shows a broad single line with slight

anisotropy similar to that observed for RufSO^F)^ The gQ value is found to be 1.724. The spectrum obtained from

Cs[Ru(SO^F)^] is quite different, and appears to contain two

overlapping lines as shown in Figure 25. The gQ values as determined from the spectrum are 1.979 and 1.84 6. It seems

that Cs[Ru(SO^F)4] contains two slightly different sites for

the paramagnetic centers with the majority having gQ value of

1.979. Structural conclusion could not be made without a crystal structure determination.

D. FLUOROSULFATO COMPLEXES OF RUTHENIUM(IV)

1. Introduction

A number of transition metals, such as silver, copper and nickel, are found with the metal in a higher oxidation state in anionic fluoro-complexes than in binary fluorides.

An analogy in fluorosulfate chemistry is reported for palladium 2- where the anion [PdCSO^F)^] is found in a number complexes 265

while Pd(S03F)4 could not be obtained . It was interesting to see whether this rationale would extend to ruthenium as well FIGURE 25 ESR Spectrum of Cs [Ru (SO-,F) A at 80 K 231

2. Preparations and Elemental Analyses

(a) K2[Ru(S03F)6)]

A 2:1 mole ratio mixture of KSO^F (0.776 g) and ruthenium

powder (0.284 g, 2.81 mmol) was loaded into a one-part pyrex

reactor inside the drybox. Approximately 5 ml of HSO^F and an

were equal volume of S20gF2 distilled successively onto the

mixture. Subsequent stirring at ^ 60 °C overnight resulted

in a red solution. The excess S~GvF_ and HSO..F were removed 2 6 2 3

in vacuo with the reactor and content kept at ^ 50 °C. Removal

of all volatiles from the viscous red-orange solution yielded

2.188 g of a bright orange solid as compared to 2.174 g expected

for K2[Ru(SO^F)g]. Result of the elemental analysis confirms the

above composition. Calculated: %Ru, 13.06; %K, 10.11; %F, 14.73.

Found: %Ru, 13.06; %K, 9.99; %F, 14.77.

(b) Cs2[Ru(S03F)6]

An identical procedure was used to synthesize Cs~[Ru (S0-.F) ,] Z. i b

from 0.255 g (2.23 mmol) of ruthenium powder and 1.034 g

(4.46 mmol) of CsS03F. The weight of the light orange solid

product was 2.152 g, compared to 2.142 g expected for

Cs2[Ru(S03F)6]. Elemental Analysis, calculated: %Ru, 10.51;

%Cs, 27.65'; %F, 11.86. Found: %Ru, 10.57; %Cs, 27.78;

%F, 11.97. 232

The reaction procedure used to synthesize M^RufSO^F)^ was adopted to the reaction of a 1:1 mole ratio mixture of

0.232 g (2.30 mmol) Ru and 0.534 g CsSC^F. The reaction was undertaken hoping to oxidize ruthenium to the +5 oxidation state and to isolate the complex Cs[RuV(SO^F)^]. Removal of all volatiles from the red solution resulted in a slightly non-homogeneous red-brown solid with dark brown impurities.

Fresh ^2^SF2 ( ^ 3 ml ) was distilled onto this initial product.

After stirring the mixture at ^ 60 °C overnight, a homogeneous red-brown powder (1.686 g) was obtained. The composition as

Cs [Ru (SO^F) j. ] was confirmed by the expected yield (1.677 g) and the elemental analysis. Calculated: %Ru, 13.86; %Cs, 18.22;

%S, 21.98; %F, 13.03. Found: %Ru, 13.68; %Cs, 18.38; %S, 22.12;

% F, 13.17.

A further attempt to oxidize Ru(IV) in Cs[Ru(SO^F)^] was made by reacting a sample Cs [Ru (SO^F) ^] with excess 820^2 at

^ 100°C. No detectable weight increase in the solid sample was observed after the volatiles were removed. It appears that ru• thenium (V) could not be obtained under the conditions described.

3. Characterizations

(a) Vibrational Spectra

Both infrared and Raman spectra were obtained on the three fluorosulfato complexes of ruthenium(IV). The Raman spectrum of Cs[Ru(SO^F)^] is rather poor due to its fairly dark red-

brown color. The observed band frequencies are listed in

Table 40. The Raman spectrum of K~[Ru(S0oF),] is illustrated

2 i b

in Figure 26.

The vibrational spectra of the hexakisfluorosulfato-

ruthenate(IV) complexes are characteristic for complexes IV 2- containing the [M (SO^F)^] ion and compare well with spectra 42 196 of K„[Sn(SO^F),] and Ba[Pt(S0oF) - ] ,listed in Table 24. 2 j b 6 b

A general assignment of the bands is made, consistent with the

presence of monodentate SO^F groups in an anionic complex.

The IR spectrum of Cs[Ru(SO^F),-] shows some resemblance to 2-

those of the [Ru^O^F)^.] spectra, except for the appearance of two bands of medium intensity at 1140 and 1045 cm 1, both

indicative of the presence of bidentate fluorosulfate groups.

It appears that each ruthenium atom in Cs[Ru(SO^F)^] is hexa-

coordinated, probably via two bridging bidentate and four mono•

dentate terminal . SO^F groups.

(b) Magnetic Susceptibility Measurements

The magnetic susceptibilities of the three ruthenium(IV)

fluorosulfato complexes were measured between 298 and 77 K.

The result is tabulated in Table 41. 4 Tetravalent ruthenium compounds with a d electronic TABLE 4 0

VIBRATIONAL FREQUENCIES OF ANIONIC FLUOROSULFATO COMPLEXES OF RUTHENIUM(IV)

K2[RU(S03F)g] Cs„ [Ru(SO-,F) ,] Cs[Ru(S03F)5] Assignment I 3 6 IR Raman IR Raman IR Raman

1400 vs, b 1405 w 1405 vs, b 1400 vw v S0o (mono) 1450 vs, b 1398 vw asy 2

v S03 (bi, A")

1261 vs 1252 vs 1252 s 1241 s 1210 vs 1215 vs v SO- (mono) 1205 vs 1216 w sym 2 1 114 0 m v S03 (>bi, A ) 1052 vs 1055 s 1057 s

1045 m v S03 (bi, A') %970 w, b ^960 w 995 mw 925 vs, b %910 vw 930 vs, b 925 vs, b v S—O (mono) 852 m ^845 vw 865 vw v S—F (mono, bi) 825 s, b 800 vw 810 s, b %820 vw 810 s, b 650 m 62 0 mw 650 mw 633 m 660 m

580 ms 589 w 580 m 580 ms S03F def. modes 55 0 ms 568 w 550 m 550 ms 460 w 4 50 mw

S03F rock, modes 440 m 44 2 mw 44 0 mw 44 0 mw 310 ms 280 m 277 m 26 9 ms M-0 str. to 270 w OJ CD (\J

OJ FIGURE 26 o10 The Raman Spectrum of

K2CRu(S03FU

* spurious band

o mOJ iOnJ ] oo . c\j

I— NJ CO 1600Ccm-n 1400 1200 1000 800 600 400 200 Ln TABLE 41

MAGNETIC SUSCEPTIBILITY DATA OF K2[Ru(S03F)g] , Cs2 [Ru(S03F) ] , Cs[Ru(S03F) ]

K2[Ru(S03F)g] Cs2[Ru(S03F) ] Cs[Ru(S03F)5]

cor m . cor „ cor _XM _ yeff T XM yeff T _% yeff T 6 3 1 6 3 1 6 3 -1 (x 10 cm mol ) (yB) (K) (x 10~ cm mol ) (yB) (K) (x 10~ cm mol ) (u ) (K)

34 01 2.85 298 3238 2 .78 298 2607 2.49 297

3654 2.83 273 3490 2 .75 271 2812 2.47 271

3957 2.81 250 3740 2.73 249 3033 2.46 249

4336 2.79 224 4083 2.70 223 3268 2.42 224

4753 2.75 200 4442 2.66 200 3600 2.39 199

5217 2.70 175 4856 2 .61 175 3938 2 .35 175

5761 2.64 152 5302 2 .54 153 4343 2.29 150

6352 2 .54 127 5800 2.44 128 4717 2.20 128

7190 2.49 108 6246 2.33 109 5149 2 .12 109

7625 2.16 76 6652 2 . 02 77 5682 1.87 77

M CO <3> configuration are expected to be low spin in an octahedral 3 environment with a T.,- g ground term. As many precedents 246 indicate , the magnetic moment should decrease with

265 temperature from % 3.0 yD at 300 K to ^ 1.7 p. at 80 K .

In particular, K^RuF^ shows a magnetic moment of 2.86 yfi at 24 6

300 K . The hexakisfluorosulfatoruthenate(IV) complexes

show a similar trends of decreasing magnetic moments from the

room temperature moments of ^ 2.8 y0 to ^ 2.1 y_ at 7 7 K. This a B is consistent with a regular octahedral coordination around

Ru(IV) as already suggested from the vibrational spectra. The

range of magnetic moment found for Cs [Ru (SO^F) is lower

(^ 2.5 yD to ^ 1.9 yD) than expected but still corresponds to 4

a low spin d ion. It appears, that antiferromagnetic coupling

perhaps by direct Ru-Ru interaction may be responsible for the

low magnetic moment. The presence of anions like [RU2(SO.JF)^Q

would still be consistent with the observation of IR absorption

bands due to bridging SO^F groups. (c) Electronic Spectra

The band maxima of the visible diffuse reflectance and mull spectra, recorded on the fluorosulfato complexes of ruthenium(IV),are listed in Table 42. Single broad absorptions 2-

are observed for the [Ru(S03F)6] complexes, with the band 3 -1 center shifting from 22.2 x 10 cm of the potassium complex 238

TABLE 4 2

ELECTRONIC SPECTRA OF FLUOROSULFATO COMPLEXES OF RUTHENIUM(IV)

AND RELATED COMPOUNDS

Compound Type of Spectrum X (x 103cm 1) max

K2[Ru(S03F)g] Mull & Diff. Refl, 22.2

Cs2[Ru(S03F)6] Mull & Diff. Refl 20.4

Cs[Ru(S03F) ] Mull & Diff. Refl 22.2 28 .6 239 Cs2 [RuFg] powder 27 .0 31.0 2- 24 5 jRuClg) solution (10 M HC1) 17.15 (600)a 20.3 (4,400) 22.9 (3,500) 24.6 (3,000) 36.0 (12,000) 41.0 (18,000)

a. extinction coefficients in M "'"cm 1 239

3 -1 to 20.4 x 10 cm for the cesium complex. This is consistent with the lighter orange-yellow color of the cesium complex. 3 -1 Besides the 22.2 x 10 cm band, an additional band at 3 -1 28.6 x 10 cm was observed in the spectra of Cs[Ru(SO^F)^]. 3 —1 2 3 9 A broad band at 27.0 x 10 cm observed for Cs,, [RuF,] has z 6 been tentatively assigned to the unresolved d—d transitions 3 between the T^ ground state and the closely spaced excited 3 3 3 3 states E , Tn , A, and A„ . On the other hand, the g 2g' lg 2g 2- observed absorptions in the solution spectrum of [RuClg]

(listed in Table 42) have been assigned to charge-transfer 245 transitions . It appears that a definite assignment of the spectra on the fluorosulfate complexes is not possible due to the rather poor resolution.

E. OTHER SYNTHETIC ATTEMPTS

1. Reaction of S„0,F„ with Ruthenium Metal 2 6 2

No apparent reaction was detected when 0.182 g of ruthenium metal was stirred in excess S2°5F2 ^ 3 m^ at room temperature.

The reaction temperature was raised to 6 0 °C for three days and subsequently to 90 °C for one day. A very slow reaction giv• ing small amounts of dark reddish brown precipitate was 240

observed. The weight of non-volatile product after reacting

for two weeks was 0.259 g. The IR spectrum of the materials

shows very weak bands in the S—O stretching region. It appears

that complete reaction was not achieved, presumably due to

surface coating that hindered further reaction.

2. Reaction of S20gF2 and Ru(S03F)3

In an attempt oxidize Ru(S0oF)_ further, excess S-0,F_

was distilled onto the product from the reaction between Ru and

HS03F/S20gF2. Stirring the mixture at 60 °C overnight and then

at 90 °C did not result in a reaction.

3. Reaction of BrS03F with Ruthenium Metal

No observable reaction was detected at room temperature

between 0.164 g (1.62 mmol) of Ru powder and excess BrSC>3F

(^ 5 ml). Stirring the mixture at 85 °C for two days resulted

in the viscous dark red solution. The volatile materials were difficult to remove in vacuo and heating of the reactor and contents to 100 °C was required to give a constant weight of

0.918 g of a black solid with traces of dark brown material.

The weight of product per mole of ruthenium was ^565 g/mol while

the molecular weight of Ru(S03F)3 is 398.25 g/mol, suggesting an impure product with little chance for a successful separation

from the excess BrS0oF. 4. Reactions of RuO

In an attempt to convert an oxide to fluorosulfate or oxyfluorosulfate, 0.095 g of anhydrous Ru"^C>2 was mixed with excess S2°gF2 ^ 3 ml) • No visible reaction was observed at room temperature and then at 60 °C. Approximately 6 ml of

HSO.jF was distilled into the one-part reactor containing the mixture. Stirring the resulting mixture at room temperature and then at 60 °C for one day each resulted in a very small increase in the weight (0.013 g) of the non-volatile material.

Excess BrSO-jF was then vacuum distilled onto the result• ing product from above, stirring the mixture at room tempera• ture for two days, at 6 0 °C for four hours, then at 100 °C for two hours resulted in a final non-volatile product weight of 0.111 g. Hence it must be concluded that anhydrous RuC^ is rather unreactive towards the fluorosulfonating agents used here. 242:

VI OSMIUM(III) FLUOROSULFATE

A. INTRODUCTION

As already mentioned, osmium and ruthenium show a similar chemical behaviour, and both differ in this respect from iron. A greater tendency for 5d block elements to exhibit higher oxidation states than 4d and 3d elements is illustrated by the fact, that the most common oxidation state of osmium is +4. The trivalent state is far less stable, and most trivalent osmium complexes are easily oxidized. However, with

TT-acceptor ligands present, reduction to the divalent state 3- 4- may also occur, as in the case of [Os(CN),] to [Os(CN),} b b As mentioned earlier, osmium octafluoride does not exist. 2 66 The synthesis of OsFg originally claimed in 1913 has been 26 7 shown to be that of OsFg . The known fluorine containing octavalent osmium compounds are Os03F2 and its complexes I I 268 M Os03F3 ( M = K, Cs, Ag ) . Osmium trioxydifluoride, is

prepared by the reaction of excess BrF3 with a 2:1 mole ratio

mixture of 0s04 and KBr. After removal of excess BrF3 at room temperature, the orange product is purified by sublimation.

Alternatively, osmium metal may be oxidized with a 2:1 volume

mixture of oxygen and fluorine to give 0sO^Fo and small amounts 243

of OsO^ and OsFg, which can be separated from the main

product (m.p. 170 °C) due to their higher volatility. The

M^OsO-jF^] complexes are made in a similar manner using a

1:1 stoichiometric mixture of MBr and OsO^.

Direct fluorination of osmium metal at 600 °C and 4 00 14 atmospheres yields OsFj , which is only stable below

- 100 °C. Base on the infrared spectrum, the structure is

suggested to be pentagonal bipyramidal. The magnetic moment measured at 90 and 195 K is approximately 1.08 u_. . Prepared 13

from the fluorination of anhydrous Os02 at 250 °C, OsOF^ is

the only other heptavalent halogen containing compound of osmium 22^. This green solid (m.p. 59.2 °C) obeys the Curie-Weiss

law with 0 = 6 K and a temperature independent moment of

1.47 y_.. The solid is orthorhombic but becomes body centered cubic above 32.5 °C 26^. The infrared spectrum of gaseous

OsOFc is consistent with a C. structure. 5 4v Using milder fluorinating condition , the yellow osmium hexafluoride is obtained from the fluorination of the metal

o 27 0 26 7 at 250 C ' . Vibrational spectra support the expected 271 octahedral symmetry . An oxychloride OsOCl^ has been prepared from the oxidation of the metal with a chlorine-oxygen mixture at 400 °C This dark brown diamagnetic solid

(m.p. 32 °C) probably contains bridging chlorines. Cs2[OsC>2Cl4] belongs to a series of "osmyl" complexes with the general formula 244 i

I I 272 I^fOsC^X^] where M = alkali metals- and X= Cl, Br, etc.

They are diamagnetic and have linear 0=0s=0 units with the

other four ligands in the equatorial plane of the compressed

octahedrons.

Like its ruthenium analogue, osmium pentafluoride contains

cis-non-linear fluorine bridging tetrameric units. It is

obtained along with osmium tetrafluoride by the reduction of 27 3 OsF, wxth W(CO), . The more volatile OsFc is vacuum DO O

distilled off the mixture at 120 °C. The yellow OsF4 melts

at 230 °C; the high melting point suggests a polymeric struc• ture. OsF5 i-sa blue-grey solid which melts at 70 °C to a green liquid and boils at 226 C to a colorless vapor '

Its magnetic moment ranges from 2.06 y0 at 295 K to 1.73 yD at

102 K. K[OsFg], along with the other alkali metal analogues, have been, prepared from the reaction of BrF^, OsBr^ and the 275 corresponding alkali metal bromide . Curie-Weiss magnetic behaviour is observed. Direct chlorination of osmium metal at 600 °C and 7 atmos- 276

pheres results in the red OsCl4 . Similarly, the correspond•

ing black OsBr4 is obtained by heating the metal with bromine

277 under pressure . A large number of hexahaloosmate(IV) complexes with the general formula M2[OsXg] is known, with 278

M = alkali metals or silver(I) and X = F, Cl, Br, or I

Their reactivity as well as magnetic and spectral properties 245

have been studied extensively and have been reviewed by 278 Griffith . In particularly, the M^OsClg] complexes have been found useful as starting materials for many reactions.-

Fluorides of osmium with oxidation state below +4 have not been reported. Osmium trichloride is best prepared by the thermal decomposition of the tetrachloride at 470 °C in a flow 276 system with a chlorine atmosphere . OsCl^ is a dark grey powder and decomposes above 450 °C to the metal. It is isomorphous with a-RuCl^- OsBr^ can be obtained from thermal 279 decomposxtion of OsBr^ . It is a black powder, and like

OsCl^, decomposes at high temperature to the metal. The corres• ponding black, amorphous Osl^ is prepared similarly by thermal 28 0 degration of (H^O^tOsI^] . Hexa-chloro-, -bromo-, and

-iodo-osmate(III) complexes have also been reported, but are not as stable as the corresponding hexahaloosmate(IV) complexes.

To summarize, the halide chemistry of osmium is very similar to that of ruthenium. The most extensively studied halo-compounds and -complexes are of osmium(IV). Fluorosulfates of osmium have not been reported. B. SYNTHETIC REACTIONS AND ELEMENTAL ANALYSES

1. Synthesis of Os(SO^F)3

In a typical preparation, excess S00cFo 3 ml) was vacuum 246

distilled onto 0.172 g ( 0.903 mmol ) of osmium powder in a one-part pyrex reactor. The resulting mixture was stirred at a constant temperature of 60 °C, and the progress of the reaction was monitored by the weight increase of non-volatile green solid. After three days, the weight of the homogeneous bright green solid reached a constant value of 0.4 31 g. The expected yield based on Os(SO^F) was 0.440 g. The elemental analysis confirms this composition. Calculated for Os(SO.jF).jJ

%Os, 39.02; %S, 19.74; %F, 11.69. Found: %Os, 39.10;

%S, 19.87; %F, 11.85, 11.88. The material decomposed above

130 °C to a black liquid.

On long standing in excess ^2^6F2 a*" room temperature over a period of several weeks, this bright green OsCSO^F)^ gradually turned into a light green solid. However, no change in sample weight was detected. In addition, the elemental analysis (%Os, 39.28; %S, 19.55; %F, 11.89 ) did not show a change in composition. Besides the difference in physical appearance from that of the bright green OslSO^F)^ (a-form), this light green sample (B-form) was found to turn to a grey color on cooling to 77 K and decompose to a black liquid at

14 0 °C. The a-form does not change color on cooling, and unlike the B-form, has very little solubility in HSO^F.

2. Other Synthetic Attempts

The reaction of osmium metal with a 1:1 by volume mixture mixture of S-0,F_ and HSO..F at ^ 65 °C yielded a viscous 2 6 2 3 brown liquid along with small amount of dark green precipitate after removal of the more volatile components in vacuo.

Homogeneous product could not be isolated.

Bromine monofluorosulfate reacted very sluggishly with osmium metal even at ^ 100 °C. Removal of excess BrSO^F after stirring at ^ 70 °C for over one month resulted in a viscous brown-green liquid and some greyish solid which appeared to be unreacted metal.

In an attempt to oxidize osmium further, excess S„0,F„ 2 o 2

was reacted with 0.440 g of the bright green (a) Os(SO.jF)3.

Stirring the mixture at ^ 70 °C for one day changed the

OsCSO^F)^ into a dark green liquid. After distilling the excess ^2^6F2 •"-n^° a storage container, the weight of the viscous dark green liquid was found to be 0.44 5 g. Further

mixing with excess S20gF2 at high temperatures up to 100 °C gave no indication of reaction. It was also found that using a reaction temperature above 60 °C when attempting to synthesize

OsCSO^F)^ would result in a similar viscous liquid.

3. Attempts to Synthesize Fluorosulfato Complexes of Osmium

Various attempts were made to synthesize fluorosulfato complexes of osmium with alkali metals. The oxidation of osmium metal by HSO^F/S^O^-F,, in the presence of different stoichiometric amounts of CsSO^F (CsSO^F/Os ratio from 1 to 3) gave non-homogeneous green-brown solids. Stepwise addition of stoichiometric amounts of CsSO-F to osmium metal in S~0,F_

only resulted in simple mixtures of Os(S03F)3 and CsSO^F, as indicated by the IR spectra of these mixtures. The addition

of S20gF2 and KSC>3F to a solution of light green (3) Os(S03F)3

in HS0oF did not oxidize osmium above +3.

C. EXPERIMENTAL RESULT AND DISCUSSION

1. Vibrational Spectra

Infrared Spectra were recorded on both forms of Os(S03F)

Raman spectra obtained were of very poor quality and showed only two bands. The band maxima are tabulated in Table 43.

The spectra of the two modifications of Os(S03F)3 appear very similar, although there are some subtle differences. Band positions are slightly shifted in the streching region. Both mono- and bi-dentate fluorosulfate groups seem to be present

in each modification of Os(S03F)3. The light green (3)

Os(S03F)3 give rise to a more complicated spectrum.

While definite assignment to the band maxima is not feasible in view of the complexity of the spectra, it appears TABLE 4 3

VIBRATIONAL SPECTRA OF Os(SO^F)

(a) (6) Raman 1235 s 1235 s

(cm-1) 1027 vs 999 vs

I.R. 1445 s 1450 s

_x 1410 s 1395 s (cm ) 1230 vs 1225 vs 1135 s,br 1160 vs,br 1075 s,br 1000 w 1025 s 950 w 950 w 890 s 880 s,br,sh 850 s,sh 800 vs 8 30 vs,br

640 m 660 m

575 s 585 s 535 s 540 m

470 vw 470 vw,sh 450 vw 450 w

395 vw 385 w 250

that the structural differences between the two forms of

OsCSO^F)^ are not apparent based on the infrared spectra,

because the relatively broad bands do not allow a clear

determination of the band maxima. Nevertheless, when compar•

ing both the relative intensities and positions of the absorptions (Figure 27), a distinction can be made between

the spectra of the two forms of OstSO^F)^.

2. Magnetic Susceptibility Measurements

Compounds of Os(III) are expected to have a low spin electronic configuration. The magnetic moment is expected to be ^ 1.9 u at room temperature and to decrease slightly with 254 temperature to ^ 1.7 u at ^ 8 0 K . The expected temperature variation of the magnetic moment is found in the limited number of variable temperature studies. However in most cases only single temperature magnetic moments are reported on 9 C osmium(III) compounds

The results of the bulk magnetic measurements on the two modifications of OstSO^F)^ are listed in Table 44. Both forms

% are weakly paramagnetic, (vef£ 0.5 yB at room temperature) and their magnetic moments decrease with temperature. The bright green (a) OstSO^F)^ shows a slightly lower magnetic moment over the temperature range studied. It appears that the low magnetic moments are best explained by antiferromagnetic 251

FIGURE 27

IR spectra of a-and £-0s(S03F)3 between 1500 and 250 cm-1

\ • i i i —i—

•

, , —,— { i 1500 WOO 1200 1000 800 GOO 400 2$0 252

TABLE 4 4

MAGNETIC SUSCEPTIBILITY DATA OF Os(SO^F)

XM yeff T XM yeff T

(x 10~6cm3mol_1) (y_) (K) (x lO^cn^mol"1) (u_) (K)

110.1 0.512 297 126.7 0.549 298

107.3 0.483 272 129.7 0.532 273

107.5 0.463 249 135.7 0.519 248

116.7 0.457 224 139.1 0.499 224

120.6 0.438 199 144.5 0.480 199

121.4 0.413 176 - 152.7 0.463 175

122.6 0.385 151 162.5 0.443 151

126.8 0.359 127 163.9 0.409 128

130.3 0.336 109 179.7 0.394 108

112.4 0.263 77 201.5 0.353 77.4 253

coupling presumably caused by direct metal-metal interactions in both modifications. A similar situation has been discussed for RuCSO^F)^. In the case of osmium(III), a more effective quenching of the paramagnetism seems to occur.

3. Electronic Spectra

Powder diffuse reflectance spectra were recorded on both

modifications of Os(S03F)3. As suggested by their difference in color, different band positions are found for the two 3 -1 modifications. The 15.4 x 10 cm band in the spectrum of

3 1 a-Os(S03F)3 shifts to 16.9 x 10 cm" for B-Os(S03F)3. For each modification, a very strong visible band is observed which

continue into the ultraviolet region. B-Os(S03F)3 dissolves

in HS03F to give a bright green solution similar in color to

that of a-Os(S03F)3. A broad weak absorption (e = 33.5) at 3-1 -2 15.6 x 10 cm was observed in a 5.07 x 10 M solution of

3-Os(S03F)3 in HS03F. Table 45 contains the electronic spectral

data obtained on Os(S03F)3 and some related compounds.

The observed absorptions in the solution spectra of 3- 3- [OsClg] and [OsBrg] have been interpreted primarily as 245 being due to charge transfer transitions , except for a possible d—d transition that gives rise to the weak band at 3-1 3- 259 22 x 10 cm in the spectrum of [OsClg] . On the other 3 -1 hand, bands observed below 40 x 10 cm in the diffuse TABLE 4 5

ELECTRONIC SPECTRA OF Os(S03F)3 AND RELATED COMPOUNDS

Compound Type of Spectrum max

a-Os(S03F) Diff. Refl 15.4 br ^25 sh >28.6 vs

6 -Os(S03F) Diff. Refl 16.9 br >28.6 vs

B-Os(S03F) HS03F solution 15.6 (33.5)'

2 ^27 (^650) (5.07 x 10" M)

3- 258 OsCl HC1 solution 22 w 32 .6 sh 35 .45 s 38 .2 s 39 .4

3- 245 OsBr, HBr solution 21 .5 (400) 25 .6 (1,400) 27 .4 (2,500) 28 .2 (3,200) 28 .8 (3,200) 31 .0 (4,000) 32 .0 (3,600) 33 .6 (4,500)

1 -1 in M~ cm 255

reflectance spectrum of K2(IrFg) have been assigned to d—d 2 6 0 2 5 9

transitions ' . The rather poor resolution and the

limited spectral range of the Os(SO^F)^ spectra do not allow a proper assignment.

4 . ESR Spectra

Broad single line spectra were obtained from a- and

3-Os (SO-.F) 0 at both 295 and 80 K. The g values determined are 3 3 o

1.997 at 295 K and 1.993 at 80 K for a-Os(S03F)3, and 2.007 at

295 K and 1.986 at 80 K for the 3 modification. These values are not significantly different, and all of them are close to

the ge for a free electron. No ESR signal was detected at

298 and 80 K from the solution of 3-Os(S03F)3 in HS.03F used

for the recording of the electronic spectrum.

Very few ESR investigations have been reported on Os(III) 92 90

systems ' . The single related study involves Na3[OsClg] 281

in HCl solution, which shows a single line at g = 1.8

It is possible that a situation similar to that of Ru(S03F)3

is involved here, but a definite conclusion cannot be made at the present moment. D. CONCLUSION

Although it seems clear that two modifications of

OsCSO^F)^ exist, differences in spectral and magnetic properties are small. It is not possible to identify the structural differences more clearly, however, it should be pointed out that the samples obtained could well contain mixtures of each modifications in different proportions.

Such a situation is encountered in RuCl^, where samples of 251 3-RuCl., often contain some a-RuCl,, VII GENERAL CONCLUSION

A. SUMMARY

The fluorosulfate chemistry of the electron-rich 4d and 5d transition metals, silver, ruthenium, and osmium, has been investigated in detail, with the emphasis placed on compounds with the metals in the higher oxidation states.

A representative number of fluorosulfate derivatives of these three elements have been synthesized in the course of this study. Structural characterizations have been based on magnetic and spectroscopic measurements. The results have been compared to those of the analogous fluorides. While the bulkier fluorosulfate groups are expected and indeed do lead to magnetically dilute systems, the anionic silver(II) complexes M^Ag(SO^F)^ are rather surprising exceptions.

AgfSO^F^ is found to form a number of derivatives, which have corresponding fluoro-analogues. Attempts to synthesize fluorosulfates of trivalent silver have been unsuccessful.

The thermal stability of the fluorosulfates may be a limiting factor, as most fluorinations of transition metals are perform• ed at fairly high temperatures (> 300 °C), temperatures 258

at which the fluorosulfate group would dissociate.

The synthetic routes ultilized on silver have applied quite well to ruthenium and osmium. The use of the oxidiz• ing mixture of S2°6F2 and HS03F nas found general applications 172 in the study of other noble metals, such as gold and 265 palladium . It is perhaps not surprising that the highest oxidation state of ruthenium obtained is +4 for the fluoro• sulfates, rather than +6 as in RuFg.

The inability to oxidize and stabilize osmium above

+3 in the form of its fluorosulfates is unpredicted, because of the relative abundance of the +4 oxidation state, and the trend of higher stablized oxidation states for the heavier

5d elements.

B. SUGGESTIONS FOR FURTHER WORK

A number of possibilities for the continuation and expansion of this study may be discussed here.

(1) The conversion of fluorosulfates to the corresponding trifluoromethylsulfates via solvolysis reactions in HSO^CF^ should extend to fluorosulfates of ruthenium and osmium.

(2) Further attempts may be made to synthesize fluorosulfates 259

of osmium and ruthenium in higher oxidation states by:

(i) Using OsO^ and RuO^ as reagents may yield fluorosulfates with the metals in oxidation states higher than +4.

(ii) The controlled direct insertion of SO^ into the M—F bonds of high valent (V,VI) binary fluorides of ruthenium and osmium may result in the corresponding fluorosulfates or more likely, mixed fluoride-fluorosulfates.

(3) Provided single crystals could be obtained, the structures of some of the compounds synthesized in this study should be determined by X-ray diffraction studies. Compounds that are

soluble in HSO-^F, such as Ru(SC>3F)3 and K2Ag(S03F)4, appear to be potential candidates.

(4) Since the presence of metal-metal interaction has been postulated for a number of compounds with unusually low magnetic moments, these compounds may be compared to some metal-metal bonded fluorosulfate derivatives synthesized

264

perhaps via the solvolysis of Ru2(C^CR)^Cl in HS03F.

(5) The detail magnetic properties of the antiferromagnetic complexes should be further investigated in order to understand the metal-metal interactions involved and also as a general search for solid state conducting properties.

(6) The synthetic methods ultilized in this study may be extended to other electron-rich 4d and 5d transition metals. 196 Some of these metals are currently under investigation 260

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APPENDIX A

LIST OF ABBREVIATIONS

vs = very strong approx. = approximately

s = strong bipy = 2,21-bipyridine

ms = medium strong py = pyridine

m - medium Me = methyl

mw = medium weak Et = ethyl

w = weak n-Pr = n-propyl

vw = very weak Diff. Refl. = diffuse reflectance

b = broad (s) = solid

sh = shoulder (1) = liquid

sym = symmetric (g) = gas *

asy = asymmetric yg = Bohr Magneton

str = stretch ESR = electron spin resonance

def = deformation UV = ultra-violet rock = rocking

IR = Infrared

R = Raman

* -20 1 Bohr Magneton = 0.92732 x 10 erg/gauss 278

APPENDIX B-l

DETERMINATION OF ELECTRON AFFINITY OF THE

FLUOROSULFATE RADICAL BY BORN-HABER CYCLE CALCULATION33 OF KSO^F

U29g KS03F + (n+l)RT + KS03F (s) K (g) S03F (g)

IP -Ea(S03F) AHj KS03F K + |^RT -|SRT

K(g) + S03F(g)

AH , K sub DS F 2 2°6 2

2" vap.S2OgF2

K(s) + |S206F2(1)

Where: E is the electron affinity of the SO,F« radical, defined as

S03F« (g) + e" S03F~ (g)

n as in MX : for KSO,F, n=l. n o

Dg Q F is the dissociation energy of gaseous S2OgF2 2 6 2 279

APPENDIX B-2

2 3 into two SO^F* radicals and determined as 97.5 kJ/mol ,

92.0 kJ/mol 23, and 91.2 kJ/mol 24.

Average D„ _ „ =93.6 kJ/mol S2°6F2

AH is the energy required to vaporize a mole of 16 2 liquid S2OgF2 = 31.9 kJ/mol

IP is the first ionization potential of K, determined K to be 4.339 eV (418.7 kJ/mol) a.

2 9 8 AH ° K is the sublimation energy of: K(s) ^ K(g) , sub a value of 90.0 kJ/mol

298 AH KS03F is the enthalpy of formation of KS03F, 35 determined to be -1155 kJ/mol

U298 KS03F is the lattice energy of KS03F at 298 K and calculated using the Kapustinskii equation,knowing the uncertain•

ties involved in assuming a NaCl type lattice for KS03F.

a. C.E. Moore, Circular of the National Bureau of Standards 467, Atomic Energy Levels Vol. Ill, p.198. b. D.J. Cubicciofti, J. Chem. Phys., 31, 1646(1959); i.b.i.d. 34, 2189, (1961). APPENDIX B-3

(418.4) 287.2 v • z+ • z 34.5 U 1 - KS03F r+ + r r+ + r_

[UT.„_ „ in kJ/mol with r and r in pm] KbU^r + —

+ C Ionic radius (Pauling) of K = 133 pm

v is the number of ions in the molecule. SO^F may be determined from the x-ray crystal 47 structure data

r d + r = (158 + 64) m = 222 m S03F- = S-F cov. F P P

or = dg^ + rcQv> Q = (143 + 66) pm = 209 pm

and may be assumed to be comparable to the Kapustinskii1s

"thermochemical radii" of isoelectronic anions.

r , - - 236 pm and r 2- = 230 pm 4 4

Using the value of 0.222 nm for r S03F

c. F.A. Cotton and G. Wilkinson. "Advan. Inorg. Chem.", 3rd Ed Wiley-Interscience, (1972) p. 52. d. F.A. Cotton and G. Wilkinson. "Advan. Inorg. Chem.", 3rd Ed Wiley-Interscience, (1972) p. 117. e. D.A. Johnson, "Some thermodynamic aspects of Inorganic Chem Cambridge, (1968), p.41. 281

APPENDIX B-4

U 611 2 kJ/m01 KS03F = •

Hence, from the Born-Haber cycle:

AH° KSO3F = - U298 KSO3F + IPK + AHgub K + § D 2 6 2

+ AH S F 1 vap 2°6 2 " 2 RT - Ea(S03F)

Ea(S03F) = - U29g KSO3F + IPK + AH^ K + \ Ds 2 b 2

+ \ AHvap S2°6F2 ~ 2 RT - AH° KSO3F

= (-611.2 + 418.7 + 90.0 + 46.8 + 15.9 - 5.0

+ 1154) kJ/mol.

E3 (SO^F) = 1110 kJ/mol (11.50 eV) a 0 ======- 282

APPENDIX C-l

DETERMINATION OF ENTHALPY OF FORMATION, AH°, OF Ag(S03F)2

BY BORN-HABER CYCLE CALCULATION

U298 Ag(S03F)2 + (n+l)RT 2+ Ag(S03F)2 (s) Ag (g) + 2 S03F (g)

-2 Ea(S03F) IPAg

AH° Ag(S03F)2 -§Sw

Ag (g) + 2 S03F (g)

DS2°6F2 AH , Ag sub ^ +AHvapS2°6F2

Ag (s) + S2OgF2 (1)

MX n = 2 for Ag(S03F)2 n

AH f Ag(S03F)2 = - U298 Ag(S03F)2 +iE2IPAg.+ AH^Ag

+ D + AH S F 3 RT 2 E (S0 F) S206F2 vap 2°6 2 " " a 3

D =93.6 kJ/mol (Appendix B) b2U6*2

AHvapS2°6F2 = 31,9 kJ/mo1 (Appendix B) 283

APPENDIX C-2

298 a AHsub Ag = 284.5 kJ/mol ; RT = 2.48 kJ/mol

iE2IP =73 kJ/mol + 2072 kJ/mol = 2803 kJ/mol

E, (SO^F) = 1110 kJ/mol (Appendix B)

The Lattice energy may be roughly estimated by Kapustinskii's

Equation; ( (418.4) 287.2 v • z • z 34 .5

U29g Ag(S03F)2 = 1 - r , + r r, + r + + (in kJ/mol)

- 2055.3 kJ/mol with r+ and r_ in pm.

Cd2+ Pd2+ r C = i^Z ± 831 = 90 pm; r_ = 222 pm (Appendix B) ;

r+ + r_ = 312 pm ; v = 3

AH° Ag(S03F)2 = -2055.3 kJ/mol + 2803 kJ/mol + 284.5 kJ/mol

+93.6 kJ/mol +31.9 kJ/mol - 3(2.48 kJ/mol)

- 2(1110 kJ/mol)

= -1070 kJ/mol

a. D.A. Johnson, "Some Thermodynamic Aspects of Inorg. Chem.", Cambridge, (1968), p.208. b. Ref. a, Appendix B. c. D.A. Johnson, "Some Thermodynamic Aspects of Inorg. Chem.", Cambridge. (1968), p.37. APPENDIX C-3

For AgF2:

AHf AgF2 = - U298 AgF2 + ^IP^ + AH^Ag + Dj

- 3 RT - 2 E (F) a

(418.4) 287.2 v • z+ • z_ 34 .5 U A F 1 - 298 ^ 2 r+ + r (in kJ/mol)

2703.2 kJ/ mol with r+ and r_ in pm.

v = 3 ; z+ = 2 z =1

r+ = 90 pm ; r_ = 136 pm

r+ + r_ = 226 pm.

D =18.9 Kcal/mol = 79.1 kJ/mol 2

34 Ea(F) = 333 kJ/mol a

d. Ref. c, Appendix B e. D.A. Johnson, "Some Thermodynamic Aspects of Inorg. Chem. Cambridge, (1968), p.43. 285

APPENDIX C-4

AH° AgF2 = -2703.2 kJ/mol + 2803 kJ/mol + 284.5 kJ/mol

+ 79.1 kJ/mol - 3(2.84 kJ/mol) - 2(333 kJ/mol)

= -210 kJ/mol APPENDIX D-l

DETERMINATION OF EFFECTIVE MAGNETIC MOMENT> u erf BY THE GOUY METHOD

The Gram or Specific Magnetic Susceptibility of the sample is determined experimentally, according to:

3 (AW - 6) W

Where 3 = Gouy tube calibration constant measured in a field

induced current of 2 amperes using HgCo(CNS)4 as a

reference.

6 = Change in weight of Gouy tube in and out of magnetic

field.

W = Weight of sample in the Gouy tube

AW = Change in weight of sample and Gouy tube in and out

of field (an average of at least six readings)

The Molar Magnetic Susceptibility:

XM = Xg (M.W.)

where M.W. = molecular weight of the sample compound. 287

APPENDIX D-2

The corrected Molar Susceptibility:

cor _ y dia XM XM iXi

d i ci

where ^iXi = sum of the diamagnetic contribution of individual

components of the sample compound (listed in

Section II .B.4) .

The effective magnetic moment:

= 2 828 r 2 ^eff -

where T = temperature (in K) of the sample.

cor

If the plot of 1/XM versus Temperature gives a linear relationship with a small intercept 0, at the temperature axis, the Curie-Weiss Law is obeyed. The Curie-

C W ma Weiss magnetic moment Veff" Y be calculated according to:

c or 2 y^W. = 2Q28 [x i (T - 0)]

where 0 = Weiss constant.