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Spectral and magnetic studies on metal fluorosulfate complexes Alleyne, Carl Stanley 1973

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c SPECTRAL AND MAGNETIC STUDIES ON METAL FLUOROSULFATE COMPLEXES by CARL STANLEY ALLEYNE B.Sc, University of B r i t i s h Columbia, 1 9 7 0 A THESIS SUBMITTED HT PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE i n the Department of Chemistry We accept t h i s thesis as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA July , 1 9 7 3 In presenting t h i s t h e s i s i n p a r t i a l f u l f i l m e n t of the requirements f o r an advanced degree at the U n i v e r s i t y of B r i t i s h Columbia, I agree that the L i b r a r y s h a l l make i t f r e e l y a v a i l a b l e f o r reference and study. I f u r t h e r agree that permission f o r extensive copying of t h i s t h e s i s f o r s c h o l a r l y purposes may be granted by the Head of my Department or by h i s r e p r e s e n t a t i v e s . I t i s understood that copying or p u b l i c a t i o n of t h i s t h e s i s f o r f i n a n c i a l gain s h a l l not be allowed without my w r i t t e n permission. Department of The U n i v e r s i t y of B r i t i s h Columbia Vancouver 8, Canada i ABSTRACT The work described i n t h i s thesis i s i n two e s s e n t i a l l y independent parts. In Chapter 3J we have studied the infrared spectra of the alkalin e earth and fi r s t - r o w t r a n s i t i o n metal bisfluorosulfates over the frequency range l+OOO - 2^0 cm i n an attempt to resolve discrepancies i n the l i t e r a -ture on the infrared spectra of these compounds. We have obtained spectra f o r these compounds which are e a s i l y assigned and we suggest that the d i s -crepancies i n the l i t e r a t u r e may be due to differences i n c r y s t a l l i n e modifications studied by the dif f e r e n t workers. We have also checked the r e l i a b i l i t y of infrared frequency s h i f t s as measures of cation-anion in t e r a c t i o n i n fluorosulfate compounds. We observe a general increase i n frequencies f o r the symmetric SO^ stretching and S-F stretching modes with increasing i n t e r a c t i o n , but no quantitative relationship i s present. However, the S-F stretching mode i s shown to give a good i n d i c a t i o n of cation-anion interaction for compounds with predominantly io n i c character. In the second part of t h i s thesis, described i n Chapter U, we have prepared tetrakis(pyridine) complexes of n i c k e l ( l l ) , copper(ll) and z i n c ( l l ) fluorosulfates. In order to determine the coordination strength of the fluoro-sulfate ion r e l a t i v e to other ions, we have also prepared and characterized complexes of the general formula Cu(py)^X 2 (X = C10^~, BF^~, N0^~, CH^CgH^SO^", CF^COO ) by infrared, electronic, and electron paramagnetic resonance spectroscopy, e l e c t r i c a l conductivity, and magnetic s u s c e p t i b i l i t y measure-ments. The r e s u l t s indicate that on the whole, the coordination strength of SO^F towards metals i s s i m i l a r , but s l i g h t l y stronger than CIO^ and BF^ , d e f i n i t e l y stronger than PF^~, and weaker than ReO,-, CH_.C,-H, SO ~ and CF_C00 . i i TABLE OF CONTENTS ABSTRACT i LIST OF TABLES v LIST OF FIGURES v i i ACKNOWLEDGEMENTS l x CHAPTER 1 : INTRODUCTION 1 1 . 1 PREVIOUS WORK ON FLUOROSULFATE COMPOUNDS 1 1 . 2 PURPOSE AND OUTLINE OF THE PRESENT WORK 1+ CHAPTER 2 : EXPERIMENTAL 6 2 . 1 APPARATUS 6 2 . 2 REAGENTS 8 2 . 3 PHYSICAL EXPERIMENTAL METHODS 1 0 2.k SYNTHESIS OF COMPOUNDS 1 £ 2 . 1 ) . . 1 General 1 5 2.K'.'2 Dehydration of Hydrated Metal Chlorides 1 5 2.lx.3 Preparation of Metal Benzoates 1 5 2.J4.I+ Preparation of Metal Bisfluorosulfates 1 6 2.U . 5 Preparation of Copper(ll) Salts 1 9 2 . ] + . 6 Preparation of Pyridine Complexes 2 0 2 . 5 ANALYTICAL DATA 2 3 CHAPTER 3 : VIBRATIONAL SPECTRA OF METAL BISFLUOROSULFATES 2 7 3 . 1 INTRODUCTION 2 7 3 . 2 EXPERIMENTAL 2 8 3 . 3 THEORY OF VIBRATIONAL SPECTRA AND BONDING 2 9 i i i 3.1+ RESULTS AND DISCUSSION 3 ^ 3.U .1 Alkaline Earth Fluorosulfates 3I4. 3 - 2 Transition Metal Fluorosulfates lj.2 3.If..3 Correlation of Infrared Band Frequencies with Cation-Anion Interaction £ 1 3'h'k Summary and Conclusions 5 9 CHAPTER hi PYRIDINE COMPLEXES OF TRANSITION METAL FLUORO-SULFATES AND RELATED COMPLEXES 6 8 U.1 INTRODUCTION 6 8 k.2 INFRARED SPECTROSCOPY 7 1 2 + . 2 . 1 Anion Bands f o r Fluorosulfate Complexes 7 1 i | . 2 . 2 Anion Bands f o r Other Complexes 7 7 1 + . 2 . 3 Pyridine Bands 8 8 k.3 ELECTRICAL CONDUCTIVITY 9 2 l+.ii ELECTRONIC SPECTROSCOPY 1 1 1 2+.U. 1 Theory of Copper(ll) Spectra 1 1 1 I|..i|.2 Results and Discussion f o r Copper(ll) Complexes 1 1 3 k'h'3 Theory of N i c k e l ( l l ) Spectra 1 2 i + h-h'h Results and Discussion f o r N i c k e l ( l l ) Complexes 1 2 6 k*5 MAGNETIC SUSCEPTIBILITY 1 3 6 Theory of Magnetic S u s c e p t i b i l i t y 1 3 6 U . 5 . 2 Results and Discussion 1 3 7 U .6 ELECTRON PARAMAGNETIC RESONANCE SPECTROSCOPY 1 L 8 i v 1+.6.1 Theoretical P r i n c i p l e s of EPR Spectra 1l|.8 J 4 . 6 . 2 Results and Discussion "I5I4. U.7 SUMMARY AND CONCLUSIONS 16 6 REFERENCES 168 APPENDIX 1 yy V LIST OF TABLES Table Page I ' I I I I I IV V VI VII VI I I IX X XI XII X I I I XIV XV XVI A n a l y t i c a l Data f o r Metal Bis f l u o r o s u l f a t e s , M(S0^F) 2 . . 2 5 Analytical" Data f o r Tetrakis(pyrMine) Complexes, M ( p y ) j X 2 . . 2 6 Correlation Table f o r and C g Point Groups 3 0 Infrared Band Assignments f o r the Alkaline Earth Fluorosulfates 3 5 Infrared Band Assignments f o r the Transition Metal Fluor o sulfates lj.3 P o l a r i z i n g Power and Infrared Band Frequencies f o r Binary Metal Fluorosulfates 5 3 Bonding and Infrared Spectral Changes i n Fluorosulfates . 61 Vib r a t i o n a l Frequency Ranges f o r Fluorosulfates 6k Assignment of Stretching Modes for Sn(SO^F)^ 6 6 Infrared Band Assignments for M(py)^(S0^F) 2 i n the So l i d State . . . 7 3 Band Assignments f o r Ni(py)^(S0^F) 2 and Cu(py)^(S0^F) 2 76 i n A c e t o n i t r i l e Solution ' Vibrations of the Perchlorate Group as a Function of Symmetry 7 8 Band Assignments f o r M(py)^(C10^) 2 i n the So l i d State . . 7 9 Band Assignments f o r Cu(py)^(BF^) 2 i n the S o l i d State . . 8 0 Band Assignments f o r Cu(py)^(Ts) 2 i n the S o l i d State . . 8 2 Vibrations of the Nitrate Group as a Function of Symmetry v i XVII XVIII XIX XX XXI XXII XXIII xxrv XXV XXVI XXVII XXVIII XXIX XXX XXXI XXXII XXXIII XXXIV XXXV XXXVI Band Assignments f o r Cu(py)^(NO^) 2 i n the So l i d State . . . 81+ Band Assignments f o r Cu(py)^(CP^C00)2 i n the S o l i d State . 8 6 Pyridine Band Assignments i n M(py)^X 2 Complexes 8 9 Pyridine In-Plane Ring Deformation ( 6 a ) for Cu(py)^X 2 Complexes 9 0 Some Properties of Non-aqueous Solvents 9 3 Conductance of Cu(py)^(S0^P) 2 i n A c e t o n i t r i l e Solution at 25°C 9 6 Onsager Slopes for Cu(py)^(S0^P) 2 i n A c e t o n i t r i l e at 25°C • 1 0 2 Molar Conductances of Tetrakis(pyridine) Complexes . . . . 1 0 5 Limiting Anion Conductances i n A c e t o n i t r i l e at 25°C . . . . 1 0 7 V i s i b l e Spectral Data for Cu(py)^X 2 Complexes 111+ Room Temperature Electronic Spectral Data f o r Ni(py)^(SO^F^ and Ni(py)^(C10^) 2 as Kel-F Mulls 1 2 7 Spectrochemical Parameters for Ni(py)^X 2 Complexes at Room Temperature 1 2 7 Crystal F i e l d Parameters f o r Ni(py)^(C10^) 2 1 3 1 Band Maxima for Ni(py)^(S0^F) 2 i n Mixed CHgClg - Pyridine Solution . . . . . . . . 1 3 3 Room Temperature Magnetic Moments of Cu(py)^X 2 Complexes • 1 3 9 Magnetic S u s c e p t i b i l i t y Data f o r Cu(py)^X 2 and Ni(py)^(SO^F) 21i4.0 Room Temperature Magnetic Moments of Ni(py)^X2 Complexes • • 11+7 EPR Parameters of Cu(py)|X 2 Complexes i n CH 2C1 2 + 1 0 % Pyridine Solution 1 6 0 Comparison of Magnetic Moments of Cu(py)^X> Complexes obtained from EPR and Magnetic S u s c e p t i b i l i t y Data . . . . 161+ O r b i t a l Reduction Factor f o r Cu(py)^X 2 Complexes 1 6 5 v i i LIST OF FIGURES Figure Page 1 Vacuum F i l t r a t i o n Apparatus . . . 7 2 Conductivity C e l l 11 3 Normal Modes of Vibration of SO^F~ ion 30 1+ Bonding and Symmetry Properties of the Fluorosulfate Group 31 5 IT- bond Overlap i n the SO^F~ Ion 32 6 Infrared Spectra of the Alkaline Earth Fluorosulfates . 36 7 The C d l 2 Structure h.0 8 Infrared Spectra of the Transition Metal Fluorosulfates 9 Suggested Structure of the Transition Metal Fluoro-sulfates 50 10 versus P o l a r i z i n g Power of the Cation 5 ^ 11 v versus Ionization Potential of the Metal 56 12 versus P o l a r i z i n g Power of the Cation 57 13 Infrared Mull Spectrum of Ni(py)^(S0^F) 2 72 11+ Infrared Spectrum of Cu(py)^(S0-^F)2 i n Dichioromethane Solution 7lj 15 Phoreogram f o r Cu(py)^(S0^F)2 i n A c e t o n i t r i l e Solution at 25°C 95 16 Shedlovsky Plot f o r Cu(py)|+(SO-^F^ i n A c e t o n i t r i l e Solution at 25°C 99 17 SLQ - - - V versus *Jl f o r Cu(py)^(S0^F)2 i n A c e t o n i t r i l e Solution at 25°C • 101 v i i i 1 8 The Relationship between the Spectroscopic Terms and the One-electron Energy Configuration f o r a Copper(ll) Ion i n an Elongated Tetragonal Octahedral Ligand F i e l d of Symmetry 1 1 2 1 9 Room Temperature Mull Spectra of Cu(py)^X 2 Complexes . • n £ 20 Solution Electronic Spectra of Cu(py)^X 2 Complexes . . . 1 1 9 g 2 1 Correlation Diagram f o r a d Configuration i n Crystal F i e l d s of 0^ and D^ Symmetry 12I+ 2 2 Room Temperature Diffuse Reflectance Spectra of N i ( p y ) u ( S 0 3 F ) 2 and H i C p y ^ C l O ^ 1 2 8 2 3 Room Temperature Kel-F Mull Spectra of Ni(py)^(S0^F) 2 and Ni(py) i +(C10^) 2 1 2 9 21+ Electronic Spectra of Ni(py)^(S0^F) 2 i n CHgC.lg and CH 2C1 2 + 25% Pyridine Solution 1 3 ^ 25 11%^ versus Temperature f o r Cu(py)^(C10^) 2 and Ni(py) i |(S0 3F) 2 • 2 6 Energy-level Diagram f o r S = 1 / 2 and I = 3 / 2 1 ^ 0 2 7 EPR Spectrum of Cu(py)^(S0 3F) ^ -'Solution at Room Temp-erature 2 8 EPR Spectrum of Cu(py)^(S0 3F) 2 Glass at Liquid Nitrogen Temperature 1 £ 6 2 9 Line Diagram of the EPR Spectrum of Cu(py)^X 2 i n Frozen Solution (without nitrogen superhyperfine s p l i t t i n g ) . . 1 £ 8 i x ACK1T0WLEDGEMEMS I would l i k e to express my gratitude to Dr. R. C. Thompson for h i s guidance and his encouragement during the course of t h i s work. Sincere thanks are due to Dr. P. Aubke with whom I had many illu m i n a t i n g discussions on the v i b r a t i o n a l spectra of fluorosulfate compounds. His readiness to help i s deeply appreciated. Thanks are also due to Dr. P. G. Herring and Dr. R. L. Tapping f o r help i n the interpretation of the EPR spectra. Acknowledgements are not complete without thanking Mr. C. T. Wong for discussions and help i n various aspects of t h i s work. 1 CHAPTER 1 INTRODUCTION 1.1 PREVIOUS WORE ON FLUOROSULFATE COMPOUNDS Fluorosulfuric acid was f i r s t described i n 1892 by Thorpe and 1 Kirman, who obtained i t by combining sulfur t r i o x i d e and anhydrous hydrogen f l u o r i d e . Since f l u o r o s u l f u r i c acid reacts v i o l e n t l y with water, they assumed that the acid i s instantaneously and completely hydrolyzed when i t comes into contact with water. Their conclusions delayed further 2 study of the acid f o r about twenty years. F i n a l l y , Traube found a method of synthesizing s a l t s of f l u o r o s u l f u r i c acid; he discovered that these s a l t s are act u a l l y f a i r l y stable i n aqueous solution, and that the acid i s incompletely hydrolyzed when i t i s dropped into water. Soon af t e r t h i s , i t was determined that perchlorate, tetrafluoroborate and fluorosulfate ions have s i m i l a r chemical properties, and that the corresponding s a l t s are very often analogous i n chemical and, to a l i m i t e d extent, i n c r y s t a l -lographic respects, and a number of mixed c r y s t a l s were prepared.3a>3^ Since that early period of investigation, f l u o r o s u l f u r i c acid and i t s derivatives have found important uses as highly a c i d i c solvent media, f l u o r i n a t i n g agents i n inorganic chemistry, catalysts, and reagents f o r ii-9 the production of important organic compounds. 2 In recent years, there has been renewed interest i n the prepara-t i o n and study of fluorosulfate compounds. This i s l a r g e l y due to the discovery that peroxydisulfuryl d i f l u o r i d e (FO^SO-OSO^F) i s an extremely v e r s a t i l e reagent i n synthetic inorganic chemistry, and i s p a r t i c u l a r l y applicable to the synthesis of a wide range of covalent fluorosulfates.^'''^''''' As a r e s u l t , the recent l i t e r a t u r e on fluorosulfate compounds i s quite extensive. Thus, there have been reports on the preparation and characterization of metal fluorosulfate compounds, including binary metal 12 , 13 v 1 i + " 1 9 + • 1 U , 2 0 , 2 1 . . , , . 21 ,22 „ ,„ . mono-, b i s - , t r i s - , ' and t e t r a k i s - ' fluorosulfates, 22 23_2(S 2T 2 8 halo- and organotin fluorosulfates, ' oxyfluorosulfates, and 29 30 complexes of t r a n s i t i o n metal fluorosulfates. Most of the work done on the coordination chemistry of fluoro-sulfates was performed i n the early part of t h i s century, when modern physical methods of characterization were not available. Lange reacted an aqueous solution containing ammonium fluorosulfate and ammonia or an amine with copper hydroxide to obtain complex copper anrmi.no fluorosulfates such as the tetrakis(pyridine) copper(ll) f l u o r o s u l f a t e , the hydrated bis(ethy-lenediamine) copper(li) f l u o r o s u l f a t e , and the anhydrous and hydrated tetraammine copper(ll) fluorosulfates. A number of other complex fluoro-sulfates were prepared by Wilke-Dorfurt and coworkers. One such series consists of the hexaantipyrine complex s a l t s of calcium, strontium, cadmium, 31 i r o n ( l l l ) , c o b a l t ( l l ) , and chromium(ill) fluorosulfates. Other complexes which have been described are the hexammine complexes of cadmium and n i c k e l ( l i ) f l u orosulfates, hexaurea chromium(lll) fluorosulfate, and the hexammine c o b a l t ( l l l ) and chromium(lll) d i n i t r a t e f l u o r o s u l f a t e . ^ 2 I t should be noted 16 that Edwards, S t i f f and Woolf found that the ammine complexes prepared i n 3 aqueous solution y i e l d sulfates rather than fluorosulfates. This throws doubt on the v a l i d i t y of much of the early work on the complexes. More 3 3 recently, Hayek, Czaloun and Krismer prepared the a c e t o n i t r i l e complexes of copper(l), antimony(lll) , and alumin.um(ipi))) fluorosulfates and studied t h e i r thermal decompositions. I t has only been within the l a s t few years that any detailed study has been made of fluorosulfate complexes with modern physical techniques. Thus, M i l n e 2 ^ prepared the t e t r a k i s ( a c e t o n i t r i l e ) complexes of copper(l), c o p p e r ( l l ) , and z i n c ( l l ) fluorosulfates and charac-ter i z e d them by means of t h e i r infrared, u l t r a v i o l e t and v i s i b l e spectra, 30 and t h e i r e l e c t r i c a l conductivity. Mayfield also prepared complexes of copper(ll) fluorosulfate with pyridine and substituted pyridine bases and characterized them primarily by magnetic s u s c e p t i b i l i t y measurements, infr a r e d , and v i s i b l e spectroscopy. I t was only a f t e r the work described i n t h i s thesis was completed, that t h i s reference was discovered. In many of the studies on fluorosulfate compounds, v i b r a t i o n a l spectroscopy played an important role i n providing information on the structure and bonding i n the compounds studied, by making use of the number and energies of the fundamental v i b r a t i o n a l modes of the fluorosulfate group. The v i b r a t i o n a l spectra of fluorosulfate compounds have also been 12—1 Jj. ii 7 used to indicate the extent of cation-anion interactions ~ •' from correlations of band frequencies with certain measures of the interactions, such as, p o l a r i z i n g power of the cation, or i o n i z a t i o n energy of the metal. The v i b r a t i o n a l spectra of binary metal fluorosulfates have thus been well studied, and those compounds whose spectra have been reported i n -1 12 13 37 d e t a i l over the frequency range l+OOO - 2$0 cm are: the a l k a l i metal, ' ' "1 Ij. "17 37 37 "iii "18 a l k a l i n e earth, ' manganese ( i l l ) , i r o n ( l l l ) , c o b a l t ( l l ) , h copperCll), 1^' 3 , 7 z i n c ( l l ) , 1 i + t i n ( l l ) , 1 ^ l e a d ( l l ) , 3 7 gallium(lll) , 2 0 21 13 21 37 g o l d ( l l l ) , s i l v e r ( l ) , platinum(lV), cadmium(ll), and mercury(ll) fluorosulfates. 1.2 PURPOSE AID 0UTLI1TE OP THE PRESENT WORK When the v i b r a t i o n a l spectrum of a fluorosulfate s a l t was studied by more than one worker, there were sometimes discrepancies i n the reported ["or 17 spectra. For example, Bernard e t . a l . observed three absorption bands for calcium fluorosulfate i n the SO^ stretching region, but Goubeau and Milne report only two bands. Goubeau and Milne also report S-F stretching vibrations f o r copper(II) and z i n c ( I I ) fluorosulfates which d i f f e r by as 1 16 37 much as 3 1 cm from values observed by other workers. Similar d i f f e r -"I ij. 37 ences i n the S-F stretching mode are also observed for the calcium, "16 18 37 and c o b a l t ( l l ) ' fluorosulfates. 0'Sullivan Mailer has reported infrared spectra f o r magnesium, calcium, barium, manganese(ll), and l e a d ( l l ) fluorosulfates which contain rather unusual band structure that cannot be e a s i l y assigned. In t h i s work we have re-examined the infrared spectra of the binary metal fluorosulfates to resolve some of the discrepancies i n the l i t e r a t u r e , and have extended the series of metal bi s f l u o r o s u l f a t e studied by infrared spectroscopy to include the i r o n ( l l ) and n i c k e l ( l l ) f l u o r o s u l -fates. We have also extended the range of known fluorosulfate compounds by preparing the tetrakis(pyridine) complexes of the n i c k e l ( l l ) , c o p per(ll), and z i n c ( l l ) s a l t s . These complexes were studied i n an attempt to determine the coordinating properties of the fluorosulfate ion r e l a t i v e to other anions. 5 Complexes of the general formula Cu(C^H^N)^X2 (X = C10^~, BF^~, CH^CgH^SO^" NO^ 5 CF^COO ) were therefore prepared and characterized i n the s o l i d state and i n solution i n order to f a c i l i t a t e the comparison. The physical tech-niques used i n the study of these complexes were infrared, electronic, and electron paramagnetic resonance spectroscopy, e l e c t r i c a l conductivity, and magnetic s u s c e p t i b i l i t y measurements. 6 CHAPTER 2 EXPERIMENTAL 2 . 1 APPARATUS  The Dry Box A D. L. Herring Corporation Dri-Lab Model HE-J43 with Dry-Train Model HE - 9 3 was used. P u r i f i e d L-grade nitrogen provided a dry i n e r t atmos-phere for the dry box. The nitrogen was circulated through the dry-train equipped with Linde molecular sieves which could be regenerated by use of an oven incorporated i n the system. The regeneration was performed p e r i o d i c a l l y to maintain a dry atmosphere. A l l of the anhydrous metal fluorosulfates and most of the pyridine complexes prepared i n t h i s work are unstable i n moist a i r . Whenever possible, a l l manipulations of the s o l i d compounds were performed i n the dry box. Once prepared, a l l of the metal fluorosulfates and the complexes were stored i n capped glass v i a l s i n the dry box or i n a desic&cator over D r i e r i t e . Fluorosulfuric Acid D i s t i l l a t i o n Apparatus Technical Grade f l u o r o s u l f u r i c acid ( A l l i e d Chemical Co.) was •30 p u r i f i e d by double d i s t i l l a t i o n i n an apparatus described previously. The main f r a c t i o n of acid b o i l i n g between 1 6 2 - 163°C was used i n the preparation of the fluorosulfates. Figure 1 - Vacuum F i l t r a t i o n Apparatus. 8 Vacuum F i l t r a t i o n Apparatus In the preparation of the compounds, a l l vacuum f i l t r a t i o n s were performed i n the apparatus shown i n Figure 1 . During t h i s procedure, when-ever i t was necessary to wash the s o l i d free of impurities with solvent, the main chamber was opened to a stream of dry nitrogen v i a stop cock A which 39 was attached to a vacuum-^ manifold. This allowed the apparatus to he f i l l e d with dry nitrogen without disconnecting i t from the vacuum l i n e . Liquid could therefore be added to wash the s o l i d without any danger of hydrolysis because the positive nitrogen pressure e f f e c t i v e l y prevents any moist a i r from entering the apparatus. 2.2 REAGENTS A l l chemicals and solvents used were Reagent or An a l y t i c a l grade unless otherwise stated. Fisher Spectranalyzed dichloromethane was used i n solutions f o r electronic spectra. Diethyl ether Diethyl ether was refluxed over calcium hydride or l i t h i u m aluminum hydride f o r several days and f r a c t i o n a l l y d i s t i l l e d . Pyridine Pyridine was refluxed over calcium hydride or barium oxide f o r two days and f r a c t i o n a l l y d i s t i l l e d . The dry l i q u i d was stored over Linde liA molecular sieves. For conductivity measurements, the pyridine was s t i r r e d f o r one day with a mixture of eerie ammonium sulfate and anhydrous 9 potassium carbonate, f i l t e r e d , refluxed over calcium hydride f o r one day, and then f r a c t i o n a l l y d i s t i l l e d . ^ Pyridine with a s p e c i f i c conductivity 8 " 1 1 of 9 x 10 ohm cm was obtained with t h i s treatment. A c e t o n i t r i l e A c e t o n i t r i l e was refluxed over calcium hydride f o r two days and f r a c t i o n a l l y d i s t i l l e d . For conductivity measurements, the solvent was refluxed over calcium hydride f o r one day and rapi d l y d i s t i l l e d . I t was then refluxed over phosphorus pentoxide f o r f i f t e e n minutes and rapi d l y d i s t i l l e d . F i n a l l y , i t was refluxed over calcium hydride f o r one day and d i s t i l l e d very slowly. This gave a c e t o n i t r i l e with a s p e c i f i c conductivity —8 —1 1 of 3 - 5 x 10 ohm cm . Dichloromethane Dichloromethane was allowed to stand over Linde i)A molecular sieves f o r several days and f r a c t i o n a l l y d i s t i l l e d . This treatment gave dichloro--11 -1 -1 methane with a s p e c i f i c conductivity of 1+ x 10 ohm cm Methanol Methanol was dried by standing over Linde 3A molecular sieves f o r several days. Nitromethane Mtromethane was p u r i f i e d three times by f r a c t i o n a l freezing i n a dry ice-acetone bath, drying over phosphorus pentoxide f o r one day, and f i n a l l y d i s t i l l i n g at 5 - 6 mm pressure.^ 2 10 Tetrabutylamroonium Bromide Bu^NBr was r e c r y s t a l l i z e d from benzene solution f o r use as a standard 1 : 1 e l e c t r o l y t e i n conductivity measurements. 2.3 PHYSICAL EXPERIMENTAL METHODS  E l e c t r i c a l Conductivity The conductivity c e l l used was based on a design by Jones and Bol l i n g e r and i s shown i n Figure 2 . The electrodes were coated with platinum black using a ch l o r o p l a t i n i c acid solution and a direct current of about 18 milliamperes f o r seven minutes, the d i r e c t i o n of the current being reversed every t h i r t y seconds to ensure equal p l a t i n i z a t i o n on both electrodes. The c e l l was steamed out and dried i n an oven at 110°C before use. The c e l l constant was determined by c a l i b r a t i n g the c e l l with aqueous potassium chloride solutions, and was about 0 . 7 U cm" . The c e l l was immersed i n an o i l bath whose temperature was regulated at 2 5 . 0 0 0 + .005°C. Measurements were taken with a precision a.c. bridge which allowed corrections for capacitances i n the c e l l . Four shunts were available i n order to measure very large resistances (up to 10 ohm) i n d i r e c t l y . A 2000 Hz o s c i l l a t o r was used as the source, and the- n u l l i n d i c a t i n g a balanced bridge was detected by the use of headphones. A l l conductivity measurements on the compounds were corrected f o r solvent conductivity. Since there was v a r i a t i o n i n the s p e c i f i c conductance of the solvent from batch to batch, the conductivity of any p a r t i c u l a r sample of solvent was always determined and the same aliquot was used for preparing the subsequent Figure 2 - Conductivity C e l l . 12 solution. Molar conductance values reported f o r the complexes represent at least three measurements of separate solutions, and the agreement i s within -1 2 -1 + 1 ohm cm mole i n a c e t o n i t r i l e solution. Magnetic S u s c e p t i b i l i t y Measurements Eff e c t i v e magnetic moments of the compounds were determined at a f i e l d strength of approximately 8 0 0 0 gauss, using a Gouy apparatus which was described previously. Measurements were made over the temperature range 315°K to 90°K. The sample tube and apparatus were calibrated using HgCo(CNS)^ as standard. For a l l the compounds studied, the res u l t s reported are the average of at least three separate determinations (each involving repacking of the Gouy tube). A l l samples were tested f o r f i e l d dependence of the magnetic moment at a f i e l d strength of approximately 1+000 gauss. The effe c t i v e magnetic moment of the metal ion, Mg£.f> was calcu-lated from the expression: where T i s the absolute temperature, and i s the molar s u s c e p t i b i l i t y corrected f o r diamagnetic s u s c e p t i b i l i t i e s and temperature-independent paramagnetic contributions. In order to determine whether the complexes obeyed the Curie law or the Curie-Weiss law, the magnetic s u s c e p t i b i l i t y data was analyzed by a least-squares treatment to obtain the Weiss constant, 6 . = 2.828-V/XTT ? e f f 13 Diamagnetic corrections were obtained from data of Eamshaw, I O -J Q Konig, and Taylor. The corrections used i n the calculation of the atomic s u s c e p t i b i l i t y of the metal ion are: Cu^ + = - 1 2 . 6 x 10 ^, N i 2 + -- 12.8 x 10~ 6, C ^ N = - U 9 . 2 x 10" 6, S0 3P" = - 1+9-2 x 10 - 6, ClO^" = - 32 x 10 - 6 , BF^~ = - 39 x 10~ 6, CH^gH^SO^" = - 97 x 10" 6. A l l diamagnetic corrections are i n c.g.s. un i t s . Infrared Spectroscopy Infrared spectra were recorded on a Perkin-Elmer 1+57 Grating Spectrophotometer covering the frequency range 1+000 - 250 cm . The c e l l windows used to contain the sample were KRS-5 (1+2% TIBr + 58% T i l ) , KBr, and AgCl. Nujol and hexachlorobutadiene were used as mulling agents. A l l mulls were prepared i n the dry box; the c e l l s were sealed with p l a s t i c tape to prevent hydrolysis, and the spectra were run as soon as possible a f t e r preparation of the mull. KRS-5 windows were used as much as possible since they allow spectra to be recorded to 250 cm ; whereas, KBr and AgCl have -1 -1 cut-off points at approximately 1+00 cm and 1+50 cm respectively. Solution infrared spectra were recorded using matched solution c e l l s with either NaCl or KBr windows, and path lengths of 0 . 1 mm, o r . 5 mm. The background absorption was always checked by running a spectrum with solvent i n both sample and reference c e l l s immediately before taking the spectrum of the sample. A l l spectra were calibrated with polystyrene f i l m at 1601.1+ cm _-| and 9 0 6 . 7 cm . 11* Electronic Spectroscopy A Cary recording spectrophotometer Model 11; was used to obtain electronic spectra i n the wavelength range 2000 - 300 nm. Solution spectra were run i n matched s i l i c a glass c e l l s with 1 or 10 cm path lengths. Solid state mull spectra were run using n u j o l , hexachlorobutadiene, or Kel-F mulls pressed between s i l i c a windows. Sample absorptions f o r nujol and hexachloro-butadiene mulls were corrected f o r l i g h t scattering as described by Taylor 18 and Thompson. Light scattering wa'S3 compensated f o r i n Kel-F mull spectra by placing a nujol-soaked f i l t e r paper i n the reference beam. Diffuse reflectance spectra were recorded on a Bausch and Lomb Spectronic 6 0 0 spectrophotometer equipped with a V i s i b l e Reflectance attachment and a Sargent recorder Model SR. Spectra were run over the wavelength range 7l+0 - 35>0 nm. Magnesium carbonate was used as the r e f l e c -tance standard. Electron Paramagnetic Resonance (EPR) Spectroscopy A Varian Associates E -3 E.P.R. Spectrometer equipped with a 100 KHz f i e l d modulation was used to record spectra at both room temperature and l i q u i d nitrogen temperature. The X-band microwave frequency was calibrated using a Hewlett-Packard Electronic Counter with a 5>256A Frequency Converter 8 - 1 8 GHz. Solutions generally had an approximate concentration of $ x 10 and were contained i n 3 mm i . d . s i l i c a tubes which were evacuated before the spectra were run. Spectrometer settings were standardized as much as possible f o r a l l the epr spectra.' 15 2.k SYNTHESIS OF COMPOUNDS  2.1;. 1 General When commercially available anhydrous metal chlorides were used i n the preparation of the metal b i s f l u o r o s u l f a t e s , the sa l t s were always flame dried and pumped on the vacuum l i n e f o r several hours to remove any surface moisture on the solids before reaction with f l u o r o s u l f u r i c acid. 2.1+.2 Dehydration of Eydrated Metal Chlorides Calcium Chloride Hydrated calcium chloride was refluxed with freshl y d i s t i l l e d thionyl chloride for two hours. The s o l i d was f i l t e r e d , and excess thionyl hq chloride was removed under vacuum. Strontium Chloride and Barium Chloride Strontium chloride hexahydrate 3 7 and barium chloride dihydrate^^ were heated i n an oven at 155°C f o r seven days. The completeness of dehydration was checked by infrared spectroscopy. 2.JU. 3 Preparation of Metal Benzoates N i c k e l ( l l ) , Copper(ll), and Z i n c ( l l ) Benzoates Concentrated solutions of sodium benzoate and the appropriate metal chloride were prepared and f i l t e r e d . They were then made s l i g h t l y a c i d i c with d i l u t e hydrochloric acid i n order to prevent the formation of basic metal s a l t s during metal benzoate p r e c i p i t a t i o n . When stoichiometric amounts of each reagent were mixed, the metal benzoate separated out of solution. The precipitate was f i l t e r e d , washed with cold d i s t i l l e d water, and dried i n a vacuum desiccator over I ^ ^ * The copper(ll) benzoate trihydrate obtained was dehydrated i n an oven at 115>°C f o r 2 days, whilst the n i c k e l ( l i ) benzoate hydrate was p a r t i a l l y dehydrated at 80°C over the same period of time. Z i n c ( l l ) benzoate was obtained anhydrous from the reaction. 18 C o b a l t ( l l ) m-Bromobenzoate Meta-bromobenzoic acid was added to a concentrated solution of sodium hydroxide u n t i l only a trace of undissolved acid remained. The solution was f i l t e r e d and allowed to react with a solution of c o b a l t ( l l ) chloride. The pink precipitate was f i l t e r e d , washed, and dried i n a desic-cator over I ^ i J * ^ e P^1^ tetrahydrated m-bromobenzoate s a l t was f i n a l l y heated to 55°C i n vacuo u n t i l i t changed to the dark blue anhydrous c o b a l t ( l l ) m-bromobenzoate. 2.I4..I4. Preparation of Metal Bisfluorosulfates In the preparation of the metal bisf l u o r o s u l f a t e s from the anhydrous metal chlorides, a l l s o l i d products were tested for the presence of chloride impurity with s i l v e r n i t r a t e solution. In a l l cases the anhydrous nature of the metal bisfluorosulfates was confirmed by the absence of absorption bands due to water i n the infrared spectra. Magnesium Fluorosulfate The compound was prepared by a modification of the procedure of 1 II Goubeau and Milne. A large excess of f l u o r o s u l f u r i c acid was d i s t i l l e d 17 onto a mixture of equimolar amounts of potassium chloride and anhydrous magnesium chloride. After the evolution of hydrogen chloride had subsided, the reaction was allowed to proceed under vacuum with magnetic s t i r r i n g f o r two weeks. Whenever the l e v e l of f l u o r o s u l f u r i c acid dropped to about h a l f of i t s o r i g i n a l volume, more acid was d i s t i l l e d into the f l a s k to bring the volume back to i t s i n i t i a l l e v e l . The s o l i d was vacuum f i l t e r e d and washed free of a l l traces of potassium fluorosulfate with f l u o r o s u l f u r i c acid. Excess acid was removed by pumping on the s o l i d product at room temperature f o r 18 hours, and f i n a l l y at 70°C for 2 hours. 37 Calcium Fluorosulfate Excess f l u o r o s u l f u r i c acid was d i s t i l l e d onto anhydrous calcium chloride and the mixture s t i r r e d at 80°C for two days. The s o l i d product was f i l t e r e d , washed several times with f l u o r o s u l f u r i c acid, and pumped free of acid on the vacuum l i n e for one day. 37 Strontium Fluorosulfate Fluorosulfuric acid was d i s t i l l e d onto anhydrous strontium chloride u n t i l a l l the s a l t dissolved. The excess acid was then pumped off on the vacuum l i n e at room temperature over a period of three days. F i n a l traces of the acid was removed by heating i n vacuo to 100°C. 37 Barium Fluorosulfate Excess f l u o r o s u l f u r i c acid was d i s t i l l e d onto anhydrous barium chloride and the mixture l e f t s t i r r i n g f o r two days at room temperature. The s a l t dissolves completely to give a clear solution. Excess acid was removed i n vacuo at 100°C over a 2{? hour period. 18 Manganese(ll), Copper(ll), Z i n c ( l l ) , and N i c k e l ( l i ) Fluorosulfates When excess f l u o r o s u l f u r i c acid was d i s t i l l e d onto the corresponding metal benzoate, and the mixture s t i r r e d magnetically, the benzoate reacted to give an immediate precipitate of the white metal fluorosulfate. In a l l cases, except f o r zinc benzoate (which was white), there was a colour change, so that the progress of the reaction could be followed v i s u a l l y . When the colour of the benzoate had disappeared completely, the s o l i d was f i l t e r e d and washed several times with f l u o r o s u l f u r i c acid. The excess acid was removed i n vacuo at 80°C over a period of one day. 1R Cobalt(II) Fluorosulfate The procedure used was the same as f o r the preceding t r a n s i t i o n metal fluorosulfates except that c o b a l t ( l l ) m-bromobenzoate was used instead of the benzoate. On reaction with acid, the colour of the s o l i d material changed from blue to pink. I r o n ( l l ) , and Nick e l ( I I ) Fluorosulfates These fluorosulfates were prepared by the procedure of Goubeau and 1 Ji Milne. Excess f l u o r o s u l f u r i c acid was d i s t i l l e d d i r e c t l y into a fl a s k equipped with a magnetic s t i r r e r and containing a mixture of equimolar amounts of the anhydrous metal dichloride and potassium chloride. The reaction mixture was protected from moisture and refluxed under a dry nitrogen atmos-phere. The r e f l u x i n g was continued f o r 7 hours f o r the i r o n ( l l ) s a l t , and 22 hours f o r the n i c k e l ( l l ) s a l t . The product was vacuum f i l t e r e d , washed with acid, and freed of excess acid under vacuum. The metal bisf l u o r o s u l f a t e s were obtained as pure white s o l i d s . 19 2.1;.5 Preparation of Copper(ll) Salts 51 Copper(ll) Nitrate Fine copper turnings ( 5 g ) were added to a solution of 1 0 g s i l v e r n i t r a t e i n 80 ml dry methanol. This suspension was l e f t f o r two days with periodic shaking. F i l t r a t i o n gave a greenish-blue solution of copper(ll) n i t r a t e which was used d i r e c t l y for reaction with pyridine to form the tetrakis(pyridine) complex. 52 Copper(ii) Trifluoroacetate Anhydrous t r i f l u o r o a c e t i c acid ( 10 ml) was dissolved i n 100 ml of absolute ethanol. Basic copper(ll) carbonate ( 2 0 g ) was added to the acid solution, and the mixture refluxed f o r three hours. This was f i l t e r e d under reduced pressure while s t i l l hot, and a dark green solution of copper(ll) tr i f l u o r o a c e t a t e was obtained which was used d i r e c t l y i n the reaction with pyridine to form the tetrakis(pyridine) complex. Copper(ll) Para-tolylsulfate Hydrate A concentrated aqueous solution of copper(ll) chloride dihydrate was added dropwise to a solution of l 5 g s i l v e r p - t o l y l s u l f a t e i n 130 ml water u n t i l no more precipitate was obtained. To determine when exactly stoichiometric amounts of each reagent were added, the reaction mixture was f i l t e r e d and a few drops of either copper(II) chloride or s i l v e r p - t o l y l -sulfate solution were added. The appearance of a precipitate indicates an excess of one reagent. When no more precipitate was obtained, the f i l t e r e d solution was concentrated on the vacuum l i n e . Hydrated copper(ll) p - t o l y l -sulfate was obtained as a pale blue s o l i d . 20 2.I4..6 Preparation of Pyridine Complexes The tetrakis(pyridine) complexes of n i c k e l ( l l ) , c o p p e r ( l l ) , and z i n c ( l l ) fluorosulfates were prepared i n the dry box by the addition of pyridine to a solution of the metal bi s f l u o r o s u l f a t e i n a c e t o n i t r i l e solution. Pure complexes were obtained only a f t e r considerable d i f f i c u l t y i n determining the proper experimental conditions. The amounts of reagents required i n the preparations were found by t r i a l and error since the purity of the complex was remarkably sensitive to the amounts of pyridine used. I f too large a volume of pyridine was used, the infrared spectrum of the complex showed evidence of l a t t i c e pyridine i n addition to the coordinated ligand. With too l i t t l e pyridine i n the synthesis, complexes were obtained which gave high a n a l y t i c a l r e s u l t s f o r metal content, and low results f o r carbon, hydrogen and nitrogen. The infrared spectra of the i r o n ( l l ) , n i c k e l ( l l ) , and z i n c ( l l ) complexes with low pyridine content showed three bands i n the asymmetric S0^ stretching region where only two are expected for unidentate coordinated fluorosulfate groups. The spectra sometimes showed bands due to coordinated a c e t o n i t r i l e . Occasionally, complexes were prepared whose infrared spectra showed three SO^ asymmetric stretching bands and the presence of l a t t i c e pyridine. These preparations evidently produced a mixture of complexes with low and high pyridine content. The experimental conditions were varied u n t i l the infrared spectrum of the complex prepared showed only two asymmetric S0^ stretching modes with no evidence of l a t t i c e pyridine, and the a n a l y t i c a l results f o r carbon i n the complex agreed with the calculated value within 1%. 21 The preparation of Fe^y^j^SO^F)^ was attempted i n t h i s work, but a l l e f f o r t s resulted i n complexes containing low pyridine and coordinated a c e t o n i t r i l e . The complex was found to be highly susceptible to a i r oxidation. B i s(fluorosulfato ) t e t r a k i s(pyridine)nickel(11) Anhydrous n i c k e l ( l l ) fluorosulfate (1.0 g) was dissolved i n 10 ml of b o i l i n g a c e t o n i t r i l e . Hot pyridine ( 6 5 ml) was added and the reaction mixture was heated to b o i l i n g . The solution was then cooled to room temp-erature and allowed to stand for t h i r t y minutes. On addition of dry ether the complex precipitated from solution. The s o l i d was then f i l t e r e d under vacuum, washed with either, and excess solvent was removed under vacuum. Bis(fluorosulfato)tetrakis(pyridine)copper(II) Anhydrous copper(ll) fluorosulfate (1.0 g) was dissolved i n 18 ml b o i l i n g a c e t o n i t r i l e . Pyridine ( 3 * 6 ml) was added and the solution boiled again. When the solution was cooled to room temperature, a purple complex immediately separated out of solution. Analysis of t h i s complex showed i t to contain pyridine i n excess of the required k : 1 mole r a t i o of pyridine to copper. The solution was allowed to stand for t h i r t y minutes and was then f i l t e r e d under vacuum. On addition of ether to the f i l t r a t e a s o l i d material precipitated from the solution. This material was f i l t e r e d from the solution, washed with ether and dried on the vacuum l i n e f o r t h i r t y minutes. Analysis of t h i s purple complex showed i t to contain the required k : 1 mole r a t i o of pyridine to copper. This complex was used for a l l measurements described i n t h i s work. 22 B i s (fluoro s u l f a t o ) t e t r a k i s(pyridine)zinc(II) Anhydrous z i n c ( l l ) fluorosulfate (1.0 g) was dissolved i n 8 ml of b o i l i n g a c e t o n i t r i l e and 1;.0 ml of hot pyridine was added. On cooling to room temperature colourless c r y s t a l s separated out of solution. The solu-t i o n was cooled i n an ice-water bath to increase the y i e l d of precipitate' produced. The c r y s t a l s were f i l t e r e d , washed with dry ether and dried under vacuum for about t h i r t y minutes. Bis(perchlorato)te t r a k i s(pyridine)copper(II) Copper(ll) perchlorate hexahydrate (5«2 g) was dissolved i n 50 ml of absolute ethanol and the solution f i l t e r e d . Pyridine ( 6 . 5 ml) was added,-and the purple complex precipitated immediately. The s o l i d was f i l t e r e d , washed with ether and excess solvent was removed under vacuum. o i B i s(tetrafluoroborato)tetraki s(pyridine)copper(II) Copper(ll) tetrafluoroborate (1+.65 g) was dissolved i n 50 ml of absolute ethanol and the solution f i l t e r e d . Pyridine ( 6 . 3 ml) was added and the complex immediately precipitated out of solution. The s o l i d was f i l t e r e d , washed with ether, and dried under vacuum at room temperature for 1-J- hours. 51 B i s(nitrato)tetrakis(pyridine)copper(II) Pyridine ( 2 5 ml) was added to a solution of copper(ll) n i t r a t e described previously (cf. section 2 . 1 ; . 5 ) . V i o l e t c r y s t a l s precipitated immediately. The solution was cooled i n an ice-water bath and f i l t e r e d under vacuum, washed with dry ether and dried under vacuum for 30 minutes. The complex was r e c r y s t a l l i z e d from hot methanol with 10% added pyridine. 23 B i s ( t r i f l u o r o a c e t a t o ) t e t r a k i s ( p y r i d i n e ) c o p p e r ( l l ) Anhydrous t r i f l u o r o a c e t i c acid ( 1 0 ml) was dissolved i n 100 ml of absolute ethanol and basic copper(ll) carbonate ( 2 2 g) was added to the solution and the mixture refluxed f o r 2-g- hours. The solution was f i l t e r e d hot and slowly added to 25 ml pyridine with thorough mixing. Blue crystals separated out immediately. The precipitate was f i l t e r e d , washed with cold ether and dried i n vacuo f o r 3 A - hours. B i s ( p - 1 o l y l s u l f a t o ) t e t r a k i s(pyridine)copper(II) Hydrated copper(ll) p - t o l y l s u l f a t e ( 2 . 7 g) was dissolved i n 22 ml of warm methanol and 8 ml pyridine was added. The complex was precipitated from solution with dry ether, f i l t e r e d under vacuum, washed with ether, and excess ether was removed under vacuum at room temperature. 80 Bis(perchlorato)tetraki s(pyridine)nickel(II) N i c k e l ( i l ) perchlorate hexahydrate ( 10 g) was dissolved i n 70 ml 2,2-dimethoxypropane and s t i r r e d magnetically f o r k hours. Pyridine ( 3 5 ml) was added, and the blue complex precipitated immediately. The s o l i d was f i l t e r e d , washed with ether and dried under vacuum at room temperature f o r k hours. 2 . 5 ANALYTICAL DATA Elemental analyses f o r carbon, hydrogen, and nitrogen were done by P. Borda of t h i s department, and su l f u r , f l u o r i n e , and barium analyses were obtained by A l f r e d Bernhardt, Mikroanalytisches Laboratorium, 5251 Elbach uber Engelskirchen, West Germany. Nick e l , copper, and zinc analyses 2h were performed by E.D.T.A. t i t r a t i o n , and i r o n ( l l ) was determined by t i t r a t i o n with potassium dichromate. The a n a l y t i c a l data are summarised i n Tables I and I I . Elemental analyses were not done on a l l the anhydrous metal fluorosulfates synthesized i n t h i s work because most of the compounds were previously characterized i n t h i s laboratory, and the same preparative procedures were followed i n each case. However, the quality of the compounds prepared were checked by infrared spectroscopy. The absence of bands due to water or sulfate i n the spectra was assurance that the preparations yielded good products. 25 Theory % Found % Metal M S F M S F Ba 1+0.91+ 1 9 . 1 2 11.33 1+0.6 1 8 . 8 1 1 . 2 Fe 2 1 . 9 8 25.21+ 11+.96 22 .0 2 5 . 1 15 .0 Ni 2 2 . 8 6 21+.97 11+.71+ — 25 .0 11+.9 Cu 21+.29 21+.51 11+.52 21+.3 21*. 6 11+.3 Zn 21+.81 21+.31+ 11+.1+2 25 .0 21+.2 11+.1+ Tahle I - An a l y t i c a l Data f o r Metal Bi s f l u o r o s u l f a t e s , M(S0-.F) Theory % Found % M X C H N M C H N M Fe SO^F" 1+2H8 3 - 5 3 9 .82 9 .79 1+1.3 3 - 3 7 9 .68 _ _ Ni 11 1+1.91 3 - 5 2 9.77 10.21+ 1+1.8 3 - 5 0 9 .50 . 1 0 . 5 Cu II 3.1+9 9 .69 10.99 1+1.2 3 . 2 7 9 .56 11.0 Zn II 1+1.1+3 3.1+8 9.66 1 1 . 2 7 1+1.0 3 - 7 0 9-72 — Ni 1+1.85 3-51 9.76 1 0 . 2 3 1+1.7 3.1+3 1 0 . 1 0 — Cu II 1+1.5o 3.1+8 9 .68 1 0 . 9 8 1+1.3 3 - 5 3 9.77 — 11 B V 1+3. ho 3.61+ 1 0 . 1 2 11.1+8 1+3.6 3 . 8 3 9 .90 — it NO^ " 1+7.67 1+.00 1 6 . 6 8 12 .61 1+7.3 3 . 9 7 1 6 . 6 7 — it p-CH-C^H,S0_" 3 6 1 + 3 5 6 . 5 3 1+.71+ 7-76 8 . 8 0 5 6 . 2 1+.91+ 7.51+ — it CF COO" 1+7.57 3 . 3 3 9 .25 10.1+9 1+7.1+ 3.1+6 9.1+7 — Table I I - Ana l y t i c a l Data for Tetrakis(pyridine) Complexes, M(py)iX, 27 CHAPTER 3 VIBRATIONAL SPECTRA OP METAL BISFLUOROSULFATES 3.1 INTRODUCTION Vibr a t i o n a l spectroscopy i s often quite useful i n obtaining information on the structure and bonding i n fluorosulfate compounds, by u t i l i z i n g the number and frequencies of the v i b r a t i o n a l bands of the fluorosulfate group. An example where i t has been used i n t h i s respect, was i n the determination of the structure of (CH-^S^ SO^F^, which was suggested to have a l i n e a r polymeric structure with bridging fluorosulfate 23 groups with the a i d of i t s Raman spectrum. This structure was subse-quently confirmed by a si n g l e - c r y s t a l x-ray d i f f r a c t i o n analysis of the compound.^ A common application of the v i b r a t i o n a l spectra of fluorosulfate 12-1L. 17 compounds i s i n i n d i c a t i n g the degree of cation-anion interactions, though there has been l i t t l e systematic work to test i t s r e l i a b i l i t y . Ruoff 12 and coworkers have studied the infrared spectra of the a l k a l i metal fluorosulfates and found that the frequency of the S-F stretching v i b r a t i o n increases with decreasing cation size i n these compounds. They concluded 28 that a s h i f t i n the position of t h i s band to higher frequency may be used as a test of increased cation-anion interaction. Bernard, Parent, and 17 Vast arrived at the same conclusion by cor r e l a t i n g the change i n the same vi b r a t i o n a l mode with the i o n i z a t i o n energy of the metals i n several binary metal fluorosulfates. In t h i s work we have studied the infrared spectra of the alkalin e earth (magnesium, calcium, strontium, barium), and f i r s t - r o w divalent t r a n s i t i o n metal (manganese, i r o n , cobalt, n i c k e l , copper, zinc) fluoro-sulfates over the frequency range 1+000 - 250 cm . This was done with two aims i n mind. F i r s t l y , to resolve some of the discrepancies i n previously reported infrared spectra, as we have mentioned i n Chapter 1, and i n so doing, to extend the series of binary metal fluorosulfates studied by infrared spectroscopy. We report here for the f i r s t time, the complete infrared spectra of i r o n ( l l ) and n i c k e l ( l l ) fluorosulfates over the indicated range. Secondly, to test the r e l i a b i l i t y of claimed trends of v i b r a t i o n a l s h i f t s with p o l a r i z i n g powers of cations or i o n i z a t i o n energies of metals as indicators of cation-anion in t e r a c t i o n . We have done t h i s by u t i l i z i n g the new data obtained i n t h i s work, and correlating a wider range of compounds than previously studied. 3 .2 EXPERIMENTAL A l l infrared spectra were recorded as described i n section 2.3. In general, reaction of the c e l l windows with the metal fluorosulfates was not observed except f o r copper(ll) and l e a d ( l l ) fluorosulfates. In the case of copper(ll) fl u o r o s u l f a t e , attack on the windows was s u f f i c i e n t l y slow that good spectra could be obtained provided they were run immediately 29 after preparation of the mull. Attack of l e a d ( l l ) fluorosulfate on c e l l windows was much more serious. I t reacts with potassium bromide plates very rap i d l y , and an infrared spectrum of potassium fluorosulfate was in v a r i a b l y obtained when these windows were used. KRS-5 and s i l v e r chloride plates were attacked less r a p i d l y , and good infrared spectra were obtained between s i l v e r chloride plates when the spectra were run immediately a f t e r mull preparation. Spectra were also obtained by shielding the plates from the mull using t h i n polyethylene sheets. Polyethylene has the advantage of being completely i n e r t to reaction with metal fluorosulfates thus eliminat-ing any p o s s i b i l i t y of the anion exchange reaction. I t s only infrared interferences i n the frequency range studied are two sharp bands at 729 cm _1 and 719 cm . 3 . 3 THEORY OF VIBRATIONAL SPECTRA AND BONDING Infrared spectroscopy i s most useful f o r characterising fluoro-sulfate compounds because of the r e l a t i v e l y high symmetry of the anion. The symmetry of the isolat e d SO^F ion i s C ^ and group theory predicts si x normal modes of v i b r a t i o n (3A^ + 3E) for t h i s species, a l l of which are infrared and Raman active. An approximate description of these modes are 53 shown i n Figure 3 - S.P. So has performed a normal coordinate analysis on the fluorosulfate anion i n the Urey-Bradley force f i e l d using the Raman frequencies obtained from solutions of sodium fluorosulfate i n water. He found that there i s some mixing between the S-F stretching and the 0-S-O and O-S-F bending coordinates. 30 vs<so5] v(S-F) V a s ( S ° 3 > 5as(S03) Figure 3 - Normal Modes of Vibration of SO^F ion I f the symmetry of the fluorosulfate group i s lowered from to C , then the degeneracy of each of the E modes i s l i f t e d and 9 normal modes s of v i b r a t i o n , a l l of which are both infrared and Raman active are expected. The co r r e l a t i o n table f o r the C-. and C point groups airs* shown i n Table I I I . Point Group A 1 A 1 A 1 E E E » i i i II i II i H C A A A A + A A + A A + A s Table I I I - Correlation Table f o r G^Y and C £ Point Groups. A reduction i n symmetry of the fluorosulfate group to C can be achieved either through s i t e symmetry lowering or through d i s t o r t i o n of the group a r i s i n g from unidentate or bidentate coordination to other atoms. I t i s usually possible to determine the source of the symmetry lowering from the magnitude of the s p l i t t i n g of the degenerate E modes i n the v i b r a t i o n a l spectra, and from s h i f t s i n band positions r e l a t i v e to those observed i n the spectra of i o n i c s a l t s such as KSO^P. s p l i t t i n g of the E modes may be observed together with a s h i f t of the band positions to higher frequencies r e l a t i v e to i o n i c fluorosulfate. The symmetries expected f o r the fluorosulfate group with different types of coordination i n t e r a c t i o n are shown i n Figure 1+. The frequency changes of the stretching vibrations may be 1 k explained by the occurrence of p7r-d7r bonding i n the ion. SO_F ion Unidentate Bidentate Tridentate Quadridentate Where there i s substantial cation-anion i n t e r a c t i o n , a large Q Q 0 ^ - 0 * C 3v C 3v C 3v Figure k - Bonding and Symmetry Properties of the Fluorosulfate Group. 32 The nature of the bonding, p a r t i c u l a r l y the importance of 3 d -o r b i t a l s i n molecules containing second-row elements such as su l f u r , i s s t i l l a matter of some speculation. However, the bonding i n the f l u o r o s u l -5i i fate ion may be r a t i o n a l i z e d i n terms of Cruxckshank's proposals on the 2 -bonding i n the isoelectronic SO^ anion. He demonstrated by structural arguments and approximate molecular o r b i t a l calculations that two 7r-bonding molecular o r b i t a l s can be formed with the 3 d x 2 y 2 and. 3&z2 o r b i t a l s of su l f u r , and appropriate 2p^ o r b i t a l s of oxygen as shown i n Figure Figure 5 - fr-bond Overlap i n the SO-F Ion. H i l l i e r and Saunders have found from ab i n i t i o molecular o r b i t a l calcu-l a t i o n s on the sulfate ion that expansion of the basis set by the addition of sulfur 3 d - o r b i t a l s r e s u l t s i n a large decrease i n the molecular energy (1 a.u.), and that the electron population of 3 d - o r b i t a l s was s i g n i f i c a n t ( 2 . 8 7 electrons). 33 As a r e s u l t of the large radius of s u l f u r , p7r-d7r overlap should be small, hence the bonding weak. However, i f the fluorosulfate group int e r a c t s , v i a coordination through oxygen atoms, with an electronegative species, such as a metal ion, then, as charge i s withdrawn by the metal, the diffuse d-orbitals are contracted, r e s u l t i n g i n better overlap between the p7r and dff-orbitals. There i s , therefore, a strengthening of the S-0 and S-F bonds due to the increased • a + ir bonding which i s r e f l e c t e d as a s h i f t by the S-0 and S-F stretching vibrations to higher frequencies with increased in t e r a c t i o n . Another factor a f f e c t i n g the v i b r a t i o n a l spectra i s the symmetry of the equilibrium configuration of the groups i n the c r y s t a l . Because of c r y s t a l influences, the free ion spectrum and s o l i d state spectrum may d i f f e r . We can discuss the symmetry of the c r y s t a l i n two parts. The f i r s t i s the symmetry of the crystallographic space group which i s concerned with the relationship among different fluorosulfate anions i n the c r y s t a l . For example, potassium and ammonium fluorosulfates belong 16 to the space group Dg^ (Pnma) with four molecules per unit c e l l . Every fluorosulfate anion i s i n the same environment as any other fluorosulfate anion, so that we need only consider one type of fluorosulfate i n i n t e r -preting the spectral data for potassium and ammonium fluorosulfates. Such may not be the case for other metal fluorosulfates, however. The other type of symmetry to be considered i s the symmetry each anion possesses because of i t s position i n the unit c e l l , i . e . , i t s s i t e 16 symmetry. In the space group , each anion i s i n one set of four 56 equivalent s i t e s of C symmetry. 3k Thus, i f the fluorosulfate anion was s u f f i c i e n t l y perturbed by i t s low s i t e symmetry i n the c r y s t a l l a t t i c e , we may see the degeneracy of the E modes removed, and the appearance of 9 v i b r a t i o n a l modes i n the infrared spectrum as expected for fluorosulfate with C g symmetry. For example, Hezel and Ross found that the spectra of the anhydrous a l k a l i metal perchlorates were consistent with the s i t e symmetry of the anion. 16 The anhydrous perchlorates they studied a3>l have the space group i n which the anion s i t e symmetry i s C . Under t h i s s i t e symmetry, a l l nine s vibrations are infrared active, and they observed a l l of these except the very weak v^. The free perchlorate ion possesses T^ symmetry which gives four fundamental frequencies, of which only two are infrared active. 3.U RESULTS AND DISCUSSION 3.1+.1 Alkaline Earth Fluorosulfates The infrared r e s u l t s and band assignments for the fluorosulfate group i n the a l k a l i n e earth fluorosulfates are given i n Table 17. The band -1 positions are considered accurate to within + 2 cm f o r sharp bands, and + 5 cm f o r broad bands. The reported frequencies are the average values obtained from several spectra, and i n most cases, several preparations. There was excellent agreement between spectra obtained from different preparations of the same compound. Figure 6 shows the infrared spectra i n _-] the frequency range 11+00 - 250 cm . Magnesium and calcium fluorosulfates show the unambiguous s i x l i n e spectra which are expected for fluorosulfate groups with G^v symmetry. 17 These r e s u l t s are to be compared with those of Bernard, Parent and Vast , 37 and 0'Sullivan Mailer, who observed rather d i f f e r e n t infrared spectra for these compounds. 35 Mg Ca Sr Ba S-0 asym. s t r . 1302 s 1270 s 1305 sb 1238 s 1313 s 1250 s S-0 sym. s t r . 111+2 s 1119 s 1113 s 1092 s 1085 s S-F s t r . 862 s 860 835 s 785 s 815 s 806 s v 5(E) S0^ asym. def. 619 m 613 s 623 m 602 m 607 m 581+ m S0^ sym. def. 571 m 568 m 586 s 5 6 5 s 567 m v,( E ) SO^ rock 1+21+ m 1+1.9 m 1+22 w 1+17 w 1+01 vw 1+15 w 1+10 w 399 w l a t t i c e mode 300 m v 2 + v3 1350 s Table IV - Infrared Band Assignments f o r Alkaline Earth Fluorosulfates 36 1400 1200 1000 800 600 400 250 Figure 6 - Infrared Spectra of the Alkaline Earth Fluorosulfates. 37 17 Bernard, Parent and Vast reported spectra f o r magnesium and -1 calcium fluorosulfates which contain two bands i n the 1000 - 1200 cm region where only one mode i s expected. The rather broad bandwidth argues against interpreting t h i s extra band as a combination or overtone band. However, the S-0 stretching vibrations a r i s i n g from sulfate groups also occur i n t h i s region. I t therefore suggests the p r o b a b i l i t y that the compounds they studied contained a s i g n i f i c a n t proportion of sulfate impurity. They also observe four bands i n the 5 0 0 - 600 cm region f o r magnesium fl u o r o s u l f a t e , and three bands f o r calcium fluorosulfate. They do not explain the excess of bands i n t h e i r spectra. 37 The spectra of 0 ' S u l l i v a n Mailer show four bands i n the 5 0 0 - 600 _1 cm region f o r magnesium fluorosulfate and three bands f o r calcium fluoro-sulfate, and a p a r t i a l l y resolved s p l i t t i n g of the band for both magnesium and calcium fluorosulfates. She r a t i o n a l i s e d that the degeneracy of the E modes i n these regions must be l i f t e d and that the extra band i n magnesium fluorosulfate must be due to a combination or overtone of l a t t i c e modes. The bands f o r both compounds are f a i r l y sharp and do not appear to be s p l i t . This band i s the most sensitive to reduction i n symmetry from to C s, and i t shows the greatest s p l i t t i n g of the three E modes. I t seems rather unusual that a s p l i t t i n g of and should be observed without any apparent s p l i t t i n g of v^. 1 ix Goubeau and Milne have also studied the infrared spectrum of calcium fluorosulfate and, as i n t h i s work, they also observe the simple 6 band spectrum f o r fluorosulfate groups. However, there are differences i n some band positions. The band frequency reported here i s some 20 cm _-| lower, and i s 18 cm higher than t h e i r reported values. -1 38 The spectrum of strontium fluorosulfate can he assigned i n terms of two non-equivalent fluorosulfate s i t e s i n the c r y s t a l l a t t i c e . I f both non-equivalent s i t e s have C^v symmetry, then we should see two s i x - l i n e spectra superimposed on each other to produce a twelve hand spectrum, provided there firss no accidental overlap of bands. In each region of the spectrum where we expect a fluorosulfate band, there i s an exact doubling of the bands expected f o r fluorosulfate. However, there are three absorption maxima i n the 11+00 - 1200 cm region where we only expect two. We may explain the band at 1350 cm as a r i s i n g from a combination of —1 —1 —1 and V ( 5 6 5 em" + 785 cm" = 1350 cm ). This combination band i s s u f f i c i e n t l y close to the fundamental mode to interact v i a Fermi resonance and gain i n i n t e n s i t y . To a f i r s t approximation, the infrared spectrum of barium fluoro-sulfate resembles one which shows fluorosulfate groups with C symmetry, s since there are nine well separated bands i n the spectrum. However, there i s also some fi n e structure on the V ^ ( E ) , v ^ A ^ ) , and V ^ (E ) bands. The _-| broad band at 121+0 cm shows some p a r t i a l s p l i t t i n g with absorption maxima -1 -1 at 1250 cm and 1230 cm . The hand also appears to be s p l i t . In _1 addition, a weak band i s observed at 1+15 cm at high mull concentrations. I f we consider a l l the fi n e structure, then we may analyse the spectrum i n terms of two non-equivalent fluorosulfate s i t e s as i n strontium fluorosulfate. Since the spectrum does not show the exact doubling of bands as ;is observed i n strontium f l u o r o s u l f a t e , there appears to be some accidental overlap of bands f o r the v H ) v_, and v i b r a t i o n a l modes. Evidently, there i s l i t t l e difference between the two proposed s i t e s i n the barium fluorosulfate l a t t i c e . The v band i s the most sensitive to changes i n i t s environment 39 as evidenced by the difference of 50 cm between both v 2 bands i n stron-_1 tium fluorosulfate, and only a difference of 21 cm f o r . The remaining bands are accordingly, less sensitive. Thus, i f we observe a difference of _1 only 9 cm between the two bands i n barium fluorosulfate, then we may not see a doubling of a l l the bands. The infrared spectrum of barium fluorosulfate has been reported 17 previously by Bernard, Parent and Vast who observed a single rather broad band, a doubled , and three-bands i n the 500 - 600 cm region. They -1 -1 -1 do not show the spectrum between 600 cm and 1000 cm or lower than 500 cm . 37 0'Sullivan Mailer shows a rather more complex spectrum with two bands and two shoulders i n the 1200 - 11+00 cm region, a single band, four bands i n the 5 0 0 - 600 cm region, and assorted unexplained bands. L i t t l e single c r y s t a l x-ray d i f f r a c t i o n work has been done on 58 fluorosulfate compounds. The only existzng studies are of KSO^P, NH^SO^P,^9 CH3C(OH)2S03P,6° (CH 3) 2Sn(S0 3F) 2, 3 6 and PXe0S0 2P. 6 1 Since no structural information i s available on the al k a l i n e earth fluorosulfates, i t i s necessary to speculate on the nature of the c r y s t a l l a t t i c e i n these compounds that i s consistent with our infrared data. Predominantly i o n i c compounds of the general formula MX2 which contain rather large cations pack so that the chief contacts are between ions of opposite sign, and so that each ion i s surrounded by the maximum number of ions of opposite sign. For moderately i o n i c compounds, such as we f i n d f o r the al k a l i n e earth fluorosulfates, t h i s may r e s u l t i n sheetlike arrangements that can be considered as approaches to spherical packings of the larger ions with the smaller ions d i s t r i b u t e d among the octahedral holes of these packings.^ 2 . In t h i s structure, a l l the octahedral holes between i+o two "close-packed" layers of X ions are occupied, and none of the holes between the next two layers. Since the primary valence of the M and X ions are s a t i s f i e d within the layer, there being only Van der Waals forces between adjacent layers, i t may be regarded as an i n f i n i t e 2-dimensional molecule. The sequence of ions are XMX XMX XMX etc., and each layer consists of octahedral MXg groups which share edges. Composite layers l i k e those shown i n Figure 7 can pack together to give cubic close-packing (ccp) or hexagonal close-packing (hep) environments f o r X. Figure 7 shows the Cdlg structure with hep of layers, which i s also found i n Calg and Mgl\-,. Such a structure f o r calcium and magnesium fluorosulfates i s consistent with t h e i r i n f r a r e d spectra. In the simple Cdl^ structure, a l l of the fluo r o s u l f a t e groups would be i n equivalent s i t e s , and we should therefore see only one type of fluorosulfate i n the inf r a r e d spectra. Figure 7 - The C d l 0 Structure. The compounds may also form mixed close-packings as i s observed for CdB^, where there i s an i r r e g u l a r r e p e t i t i o n of the cubic and hexagonal close-packings. I f t h i s i s the case f o r strontium and barium fluorosulfates, then we should detect the presence of non-equivalent anion s i t e s from t h e i r infrared spectra. A mixture of ccp and hep i n the c r y s t a l s of magnesium, calcium and barium fluorosulfates w i l l explain the infrared spectra observed 17 37 f o r these compounds by Bernard, Parent and Vast and 0'Sullivan Mailer. Different preparative procedures used by each worker may give products which d i f f e r i n the r e l a t i v e amounts of each type of packing i n the c r y s t a l l a t t i c e . Thus d i f f e r e n t c r y s t a l l i n e modifications of the same compound may have been studied by these workers thereby giving infrared spectra which are not consistent with each other. For example, cadmium hydroxide can be obtained with variable hep or ccp i n the l a t t i c e simply by changing the rate of go p r e c i p i t a t i o n of the compound from solution. The structures of fluorosulfate s a l t s have often been compared with those of the corresponding perchlorates, tetrafluoroborates, and Six 13 permanganates. Thus, Lange and Sharp have shown from gross morphology and powder x-ray d i f f r a c t i o n studies that the a l k a l i metal fluorosulfates are usually isomorphous with the corresponding perchlorates. However, there does not seem to be any s i m i l a r resemblance f o r the a l k a l i n e earth s a l t s . 65 DePape and Eavez observed from powder x-ray d i f f r a c t i o n studies that the a l k a l i n e earth tetrafluoroborates are not isomorphous with the corresponding perchlorates and permanganates. They also found that calcium tetrafluoro-borate i s isomorphous with the strontium s a l t , i n contrast to the corresponding fluorosulfates whose infrared spectra indicate no s i m i l a r relationship. 42 Although comparisons of the structures of the alkal i n e earth s a l t s are of l i m i t e d use, i t may s t i l l be i n s t r u c t i v e to consider t h e i r structures. However, the only report on the structure of any compound of the type M(AX^)2 seems to be a powder x-ray d i f f r a c t i o n study of barium . 66 permanganate. Barium permanganate c r y s t a l l i z e s i n the orthorhombic space group I>2h (Fddd) with eight formula" units i n the c e l l . The barium atoms have been determined at the center of an orthorhombic prism of eight oxygens. 9) The s i t e symmetry of the barium atoms i s D^. For the Bv,^  space group, there are also three non-equivalent s i t e s with point symmetry which may be f i l l e d with the permanganate groups. Of these three s i t e s , only one s i t e i s found to be f i l l e d with the permanganate groups. 3 . 4 ' 2 Transition Metal Fluorosulfates The infrared r e s u l t s and band assignments for the t r a n s i t i o n metal and l e a d ( l l ) fluorosulfates are l i s t e d i n Table V, and the spectra are shown i n Figure 8 . The spectra shown for the fluorosulfates were obtained as nujol mulls between KRS-5 plates except f o r l e a d ( l l ) fluorosulfate which i s shown between s i l v e r chloride plates, and copper(ll) fluorosulfate which was run between potassium bromide plates. Spectra were also obtained with very concentrated mulls i n order to detect any weak bands i n the low frequency region. Copper(ll) fluoro-sulfate shows a weak l a t t i c e mode at normal mull concentrations. Figure 8 shows t h i s band for a thick mull. The spectrum of l e a d ( l l ) fluorosulfate was obtained before any s i g n i f i c a n t attack of the c e l l windows had taken place. This spectrum i s i d e n t i c a l i n appearance and band positions with spectra obtained from mulls Mn(S0 3P) 2 Pe(S0 3F) 2 Co(S0 3F) 2 Ni(SO P ) g Cu(S0 3F) 2 Zn(S0 3F) 2 Pb(S0 3F) 2 T ^ ( E ) 1297 s 1261 s 1266 s 1262 s 1306 s 1266 s 1265 sb S-0 asym. s t r . 1223 s 1219 sh ^(A.,) 1118 m 1118 s 1111+ s 1120 s 1115 s 1118 s 1071+ s S-0 sym. s t r . V2(A.,) 836 s 862 m 858 m 859 m 861 m 859 m S-F s t r . 826 m V ^ ( E ) 601+ m 610 s 611+ m 619 m 63O m - 613 s 599 m S-0 asym. def. 590 m 601+ m 577 m .p-V3(A.,) 573 m 568 m 568 m 568 m 561+ m 567 m 559 m S-0 sym. def. T>6(E) 1+18 w 1+19 m 1+22 m 1+22 m 1+28 m 1+19 m 1+08 w S-F def. 1+12 sh 1+16 m 399 w l a t t i c e mode 290 m 300 m combination 1359 m 970 w or overtone 568 msh 760 m Table V - Infrared spectra of t r a n s i t i o n metal fluorosulfates (cm ) 1400 ' 1200 ' 1000 ' 800 ' 600 ' 400 250 Wavenumber (cm ) Figure 8 - Infrared Spectra of the Transition Metal Fluorosulfates. 45 1400 1200 1000 800 600 400 250 Wavenumber (cm ) Figure 8 continued. 46 pressed between polyethylene films and c e l l windows. The Vg bands which are obscured by the s i l v e r chloride absorption were obtained with KRS-5 plates and polyethylene f i l m s . We can consider the spectra of the t r a n s i t i o n metal fluorosulfates i n two groups. One group shows the simple spectra f o r the fluorosulfate anion with point symmetry. The other group of compounds contains more bands i n the infrared spectra than we would expect for C^v fluorosulfate anions. The infrared spectra of i r o n ( l l ) , c o b a l t ( l l ) , n i c k e l ( l l ) , and zi n c ( I I ) fluorosulfates show unambiguous s i x l i n e spectra t y p i c a l of the fluorosulfate group with symmetry. A l l bands are r e l a t i v e l y sharp, and there i s remarkably l i t t l e difference i n band positions between them. Of these four fluorosulfates, only the spectra of c o b a l t ( l l ) and z i n c ( l l ) fluorosulfates have been reported i n d e t a i l . The band positions obtained for c o b a l t ( l l ) fluorosulfate show very good agreement with the values of Taylor *I 8 "1 and Thompson, the largest difference being 8 cm i n the mode. On the other hand, the agreement of z i n c ( l l ) fluorosulfate with the work of Goubeau 14 1 and Milne i s not quite as close, there being a difference of 27 cm i n the v i b r a t i o n a l frequency. A closer agreement within 9 cm f o r t h i s -1 1 6 band i s obtained with Woolf's reported value of 850 cm . Goubeau and 14 —1 Milne observed a weak l a t t i c e mode at 261 cm which was not observed here. This i s very close to the low frequency l i m i t of the infrared spectro-photometer, and the band i s most probably hidden by the absorption of the KRS-5 plates used. The fluorosulfate s a l t s i n the second group are copper(II), manganese(ll) and l e a d ( l l ) fluorosulfates. The spectrum of copper(ll) fluorosulfate c l e a r l y shows that the symmetry of the fluorosulfate anion i s reduced to C . The degeneracy of the E modes are removed, and we s 1+7 therefore observe a s p l i t t i n g of these bands to produce an overall nine band spectrum. The infrared spectrum of copper(ll) fluorosulfate has been "1 l i 37 reported previously by Goubeau and Milne, and 0 ' S u l l i v a n Mailer. There 1 li i s f a i r agreement with the values of Goubeau and Milne, the greatest difference being i n the stretching modes. The frequency obtained here i s higher than t h e i r value by some 19 cm . We did not observe any l a t t i c e mode at 262 cm as they report, possibly due to KRS-5 absorption. We obtain good agreement with the frequencies of 0 ' S u l l i v a n Mailer, a l l _1 frequencies agreeing within 10 cm The assignment of the: :infrared spectrum for manganese(ll) fluoro-sulfate i s somewhat more ambiguous. I t shows neither the simple s i x bands for C_ fluo r o s u l f a t e , nor the nine bands for C fluorosulfate. In the 3 v s -1 -1 1+10 - 1+20 cm region, there i s a weak band at 1+18 cm and a shoulderaat _-| 1+12 cm which i s not very evident i n Figure 8 . However, by using a concen-trated mull, slow scan speed, and narrow s l i t width, t h i s can be resolved into two separate bands. These two bands are assigned to the two components of when the degeneracy of the E mode i s l i f t e d . In the 550 - 620 cm region, there are three bands which are expected f o r the s p l i t band? _-| and Vy A shoulder at 568 cm i s evident on the band, and t h i s may _-] be due to a combination of overtones of l a t t i c e modes. In the 1200 - 11+00 cm region, we would also expect the remaining E mode to be s p l i t to give effe c t i v e symmetry f o r the fluorosulfate group. There i s a sharp band -1 -1 at 1359 cm and another strong, broader band at 1297 cm . I f the band was s p l i t , the two bands produced ought to be of simi l a r i n t e n s i t y as i s 37 observed f o r copper(ll) fluorosulfate (Figure 8 ) , and li t h i u m fluorosulfate. No example i s known for covalently bonded unidentate, bridging bidentate or tridentate fluorosulfates where there i s a great i n t e n s i t y difference between 48 the two components of i n the infrared spectrum. I t i s more reasonable to assign the 1359 cm band as a r i s i n g from an overtone or combination of fundamental modes whose i n t e n s i t y i s enhanced by Fermi resonance with the nearby band. I t appears that there i s only a very small s p l i t t i n g of which cannot be resolved i n the infrared spectrum due to the natural width of t h i s band. Attempts to obtain the Raman spectrum of t h i s compound i n order to observe t h i s s p l i t t i n g were unsuccessful. The infrared spectrum of manganese(ll) fluorosulfate therefore shows a small reduction i n symmetry from to C g f o r the fluorosulfate group. The appearance of the infrared spectrum of l e a d ( l l ) fluorosulfate i s s i g n i f i c a n t l y d i f f e r e n t from those of the other metal fluorosulfates. The band i s broader than those of the other s a l t s studied, and i t -1 -1 shows a maximum absorption at 1265 cm with a shoulder at 1219 cm , ind i c a t i v e of a s p l i t t i n g of the E modes. The two components of t h i s band overlap to give the band shape observed. The remaining E modes are also -1 -1 seen to be s p l i t . There are two bands at 599 cm and 577 cm f o r -1 -1 and a broad band at 399 cm with a shoulder at i+08 cm f o r Vg. There are -1 -1 also additional bands at 7^0 cm and 970 cm which may be assigned to combination and overtones of the fundamental modes. The infrared spectrum of l e a d ( l l ) fluorosulfate has been reported previously by 0 ' S u l l i v a n 37 Mailer, who observed a rather d i f f e r e n t spectrum with di f f e r e n t band shapes, and an abundance of structure. Due to the m u l t i p l i c i t y of structure, she was unable to assign a l l the bands. For t h i s reason, the spectrum of l e a d ( l l ) fluorosulfate was reinvestigated, and the spectrum obtained shows a much simpler appearance than that previously reported. When the nujol _1 mull i s attacked by the c e l l windows, the combination band at 970 cm _-| disappears and the band at 826 cm decreases i n i n t e n s i t y while a new 1+9 band appears at about 730 cm . The spectrum of l e a d ( l l ) fluorosulfate 37 reported previously probably arises from some attack of the compound with the c e l l windows, since l e a d ( l l ) fluorosulfate was found to r e a d i l y attack KRS-5> which was used i n the previous work. The infrared spectrum of l e a d ( l l ) fluorosulfate i s sim i l a r to that of t i n ( l l ) f luorosulfate. B i r c h a l l and coworkers observed two bands -1 -1 at 830 cm and 773 cm which they assigned to the 2Vg overtone, and the Vg fundamental mode respectively. In t h i s work on l e a d ( l l ) fluorosulfate, we assign the higher frequency band to v^, and the band at 76O cm to an overtone or combination of normal modes. In t h i s respect, we disagree with t h e i r assignment of t i n ( l l ) f l u o rosulfate. We were unsuccessful i n obtaining the Raman spectrum of l e a d ( l l ) fluorosulfate. Prom infrared and magnetic s u s c e p t i b i l i t y data, Taylor and 18 Thompson concluded that there i s octahedral coordination of the cobalt ions i n c o b a l t ( l l ) fluorosulfate, r e s u l t i n g i n a CoOg skeleton which i s not s i g n i f i c a n t l y distorted from 0 ^ symmetry. In i t s more i o n i c s a l t s such as those of the a l k a l i metals, where cation-anion i n t e r a c t i o n i s r e l a t i v e l y weak, the fluorosulfate ion retains C-. character. However, i n the t r a n s i -3 v t i o n metal fluorosulfates, cation-anion i n t e r a c t i o n should be s u f f i c i e n t l y important to a l t e r the v i b r a t i o n a l frequencies of the fluorosulfate group. For these compounds we observe a s i g n i f i c a n t s h i f t of the fundamental modes to higher frequencies, especially i n v^. Since the fluorosulfate ion s t i l l maintains symmetry f o r the i r o n ( l l ) , c o b a l t ( l l ) , n i c k e l ( l l ) and z i n c ( l l ) f luorosulfates, i n the presence of s i g n i f i c a n t cation-anion i n t e r -action, a l l three oxygen atoms should be coordinated i n an equivalent manner to the metal ions. The fluorosulfate ion therefore bonds i n a tridentate fashion with the t r a n s i t i o n metal ions. The polymeric l a t t i c e proposed f o r the t r a n s i t i o n metal fluorosulfates i s shown i n Figure 9« 50 Figure 9 - Suggested Structure of the Transition Metal Fluorosulfates. 51 The infrared spectrum of copper(ll) fluorosulfate shows that the symmetry of the fluorosulfate ions are reduced from to Cg. As others have pointed out previously, the symmetry of the CuOg octahedron i s probably with four of the oxygen atoms i n a square plane closer to the copper ion than the two oxygen atoms above and below the plane. The two apic a l oxygen atoms therefore interact more with the copper ion than the a x i a l oxygens, r e s u l t i n g i n the reduction i n symmetry to C . A simi l a r s d i s t o r t i o n of the MnOg octahedron may also occur i n manganese(ll) fluoro-sulfate . 3.1+.3 Correlation of Infrared Band Frequencies with Cation-Anion Interaction As we have noted i n chapter 1, v i b r a t i o n a l spectra have often been used to indicate the extent of cation-anion interactions i n the binary metal fluorosulfates. In t h i s respect, a s h i f t to higher frequency of and base usually been taken to determine the strength of t h i s - i n t e r a c t i o n . The only systematic test of these assumptions ha-se been performed on the a l k a l i metal fluorosulfates and some alkalin e earth and t r a n s i t i o n metal 12 17 fluorosulfates. ' In t h i s section, we examine the available infrared data to see i f there i s any correla t i o n of band positions with some property of the metal i n order to determine the presence of any general relationship between the band positions and the extent of cation-anion interaction. In the s a l t s studied, when the cation and anion are juxtaposed, there w i l l be an in t e r a c t i o n over and above t h e i r coulombie a t t r a c t i o n , namely an in t e r a c t i o n due to mutual po l a r i z a t i o n . Since cations are usually small and compact r e l a t i v e to the anions, the po l a r i z a t i o n of the cation by the anion may be considered to be r e l a t i v e l y n e g l i g i b l e . Ve need, therefore, 52 only consider p o l a r i z a t i o n by the cation of the anion, which i n t h i s case, i s the fluorosulfate anion. The p o l a r i z i n g power of the cation depends upon both the i o n i c potential of the cation and on how e f f e c t i v e l y the cation nucleus i s shielded from the anion by the electrons of the cation. A quantitative measure of the p o l a r i z i n g power of the cation has been given 67 by Janz and James: P = Z 5 z1 ' 2 7 r l where P = p o l a r i z i n g power, Z = i o n i c charge, r = i o n i c radius(jl) and I = i o n i z a t i o n potential (eV). One should be aware however, that from the nature of the wave functions f o r atoms and ions, they have no d e f i n i t e s i z e , and we cannot specify the exact position of the electron clouds. There are several methods available f o r estimating the i o n i c r a d i i , the most well known being those of Pauling and Goldschmidt. Several workers have revised t h e i r values, and the most recent and comprehensive set of r a d i i f o r metal ions has been compiled by Shannon and P r e w i t t . ^ They obtained t h e i r values empirically by c o r r e l a t i n g approximately one thousand interatomic distances i n metal oxide and metal flu o r i d e structures. Adjustments were made for factors such as high- or low-spin state of the metal ion and v a r i a t i o n i n the coordination number of both cation and anion. The values of Shannon and Prewitt's^effective i o n i c r a d i i represent the best available estimate of i o n i c sizes. 53 Cation l(eV) ( 6 9 ) r(I) (68) P v 2(cm" 1) V 1(cm - 1) Reference L i + 5 . 3 9 0 . 7 U 1.51 812 1112 37 Na + 5 . 1 3 8 1 . 0 2 0 . 9 6 775 1095 37 K+ k.339 1 . 3 8 0 . 6 6 71+9 1079 37 Rb + 1+.176 1.1+9 0 . 6 0 729 1076 37 Cs + 3 . 8 9 3 1 . 7 0 0 . 5 1 716 1071+ 37 - - 0 701+ 1068 37 Mg 2 + 2 2 . 6 7 5 O.72 2 . 3 9 862 111+2 t h i s work C a 2 + 1 7 . 9 7 9 1 . 0 0 1.61+ 860 1119 » » S r 2 + 1 6 . 9 8 9 1 .16 1-35 8 3 5 , 785 1 1 1 3 , 1092 „ B a 2 + 1 5 . 2 1 1 . 3 6 1 . 1 2 8 1 5 , 806 1085 Mn 2 + 2 3 . 0 6 8 0 . 8 3 0 1.91 836 1118 F e 2 + 2 U . 0 5 0 . 7 8 0 2 . 0 6 862 1118 C o 2 + 21+.91 0.71+5 2 . 1 6 858 1111+ M 2 + 2 5 . 7 8 0 . 6 9 0 2 . 3 9 859 1120 » C u 2 + 2 8 . 0 1 0 . 7 3 2 . 1 0 861 1115 Z n 2 + 2 7 . 3 5 0 . 7 5 0 2.01+ 859 1118 „ „ P b 2 + 2 1 . 9 7 0 1 . 18 1 .16 826 1071+ S n 2 + 22.14+3 1 . 2 2 1 . 0 9 830 1062 15 C d 2 + 2 5 . 8 9 5 0 . 9 5 1.1+7 855 1107 37 Hg 2 9 . 1 8 1 . 0 2 1 . 2 5 856 1089 37 F e 3 + 51+.69 0 . 6 5 3.1+8 850 1090 11+ G a 3 + 5 7 . 2 7 0 . 6 2 3 . 6 5 868 1090 20 Table VI - P o l a r i z i n g Power and Infrared Band Frequencies for Binary Metal Fluorosulfates. POLARIZING POWER OP METAL Figure 10 - V 2 ™ - P o l a r i z ^ f f Power of Metal. 55 Table VI l i s t s the p o l a r i z i n g power of the cations and the i n f r a -red data used i n obtaining Figures 10 and 1 1 . The i o n i z a t i o n potentials l i s t e d f o r the mono-, d i - , and t r i v a l e n t metal cations are f o r the f i r s t , second, and t h i r d i o n i z a t i o n potentials respectively. In Figure 1 0 , the band frequencies f o r the metal fluorosulfates are plotted against the p o l a r i -zing powers of the cations. The graph shows an i n i t i a l trend toward higher U>2 frequencies with increasing p o l a r i z i n g power, P, of the cation, but at higher values of P, the Vg band becomes quite insensitive to changes i n P. Since there i s considerable scatter of points on the graph, care should be taken i n using the band position as a quantitative indicator of the extent of cation-anion in t e r a c t i o n i n the fluorosulfates. While there i s a f a i r l y good correl a t i o n f o r the more io n i c fluorosulfates, t h i s trend i s much more tenuous i n the more covalent fluorosulfates, and use of the frequency i s l i k e l y to lead to the wrong conclusions. For example, i n the infrared spectra of i r o n ( l l ) and i r o n ( l l l ) fluorosulfates, the band positions -1 -1 are 862 cm and 850 cm respectively, while the p o l a r i z i n g powers of the cations are 2.11+. and 3 « 5 6 . The band frequency i s actually lowered by 12 cm f o r an increase i n p o l a r i z i n g power of 1.1+2! Whi]e we may use to disti n g u i s h between i o n i c and the more covalent fluorosulfates, i t i s of l i t t l e use when the fluorosulfates being compared have predominantly covalent character. 17 Bernard, Parent, and Vast have studied the changes i n the band frequencies as a function of the i o n i z a t i o n energies of the metals, and conclude that there i s an approximate straight l i n e relationship. A re-examination of the relationship with the band frequencies obtained recently (Table VI) i s shown i n Figure 1 1 . There i s a good l i n e a r r e l a t i o n -ship between v0 and the f i r s t i o n i z a t i o n potentials of the a l k a l i metal Ionization Potential (eT), Figure 11 - V ? vs. Ionizing Potential of Cation. 58 fluorosulfates. The second i o n i z a t i o n potentials of the divalent metals do not show any s i m i l a r correlation and there i s a large scatter of points i n the graph. As we observe i n Figure 1 0 , the plot of i o n i z a t i o n potential as a function.of the frequency shows a large i n i t i a l increase i n the frequency, then a l e v e l l i n g o f f at high values of i o n i z a t i o n p o t e n t i a l . Figure 11 again reinforces the contention that Vg band frequencies may only be useful as a measure of the extent of cation-anion i n t e r a c t i o n f o r the more i o n i c fluorosulfates, or when i t i s used to compare the covalent character i n two metal fluorosulfates, i t i s only r e l i a b l e when there i s a f a i r l y large difference i n cation-anion i n t e r a c t i o n between the fluorosulfates. The v.| band frequency has also been used to measure the extent of cation-anion i n t e r a c t i o n , but t h i s has not been used often, since the band pos i t i o n i s less sensitive to changes than the band. Figure 12 shows the band frequencies of anhydrous metal fluorosulfates plotted as a function of the p o l a r i z i n g power of the cation using data taken from Table VI. There i s a general trend to higher frequencies f o r with increasing p o l a r i z i n g power of the cation. Again because of the large scatter of points, the band i s not a good quantitative measure of cation-anion i n t e r a c t i o n . For example, the infrared band frequencies for t i n ( l l ) f l u o r o s u l f a t e , and tetraphenylarsonium fluorosulfate are -1 -1 1062 cm and 1068 cm , while t h e i r p o l a r i z i n g powers are taken as 1 . 0 9 and 0 respectively. This i s the opposite order than we would expect •>om from t h e i r p o l a r i z i n g powers. 59 3'k'h Summary and Conclusions In t h i s chapter we have attempted to resolve some of the discrepancies i n the l i t e r a t u r e on the infrared spectra of a l k a l i n e earth fluorosulfates. We have obtained infrared spectra of Mg(SO^F^ and Ca(S0 3F) 2 which show simple 6 band structure assignable to C ^ f l u o r o s u l -1L fate groups. Goubeau and Milne also observed a 6 band spectrum for "17 37 Ca(SC>3F)2, but Bernard et a l . and 0'Sullivan Mailer reported rather complex spectra f o r the magnesium and calcium fluorosulfates, with structure i n excess of the 6 bands observed here. We suggest that the discrepancies between a l l the reported infrared spectra may be due to differences i n c r y s t a l l i n e modifications of the compounds studied by the different workers. Thus, MgCSO^F^ and Ca(S0 3F) 2 prepared i n t h i s work, and Ca(S0 3F) 2 prepared by Goubeau and 1 II Milne contain fluorosulfate groups i n only one type of anion s i t e i n the 17 c r y s t a l l a t t i c e . In comparison, the preparations of Bernard et a l . and 37 0'Sullivan Mailer yielded compounds with fluorosulfate groups i n two non-equivalent anion s i t e s within the c r y s t a l l a t t i c e . In addition, there 17 i s the p o s s i b i l i t y that the compounds studied by Bernard et a l . contained s i g n i f i c a n t amounts of sulfate impurity. The infrared spectra of S r ( S 0 3 F ) 2 and Ba(S0 3F) 2 obtained here show direct evidence of the presence of non-equivalent fluorosulfate s i t e s i n the l a t t i c e , though the spectrum i s not as well resolved for,Ba(S0 3F) 2 as f o r Sr(S0 3F) 2. The infrared spectrum of Fb(SC>3F)2 obtained here i s simpler i n 37 appearance to the one reported by 0'Sullivan Mailer, 1 and i t can be assigned i n terms of C symmetry for S0 QF~, with two additional combination or over-37 tone bands. The spectrum observed previously probably arises from the 60 anion exchange reaction of Ft^SO^F^ with the c e l l windows. The t r a n s i t i o n metal fluorosulfates studied show infrared spectra which are e a s i l y assigned on the basis of symmetry f o r the fluorosulfate groups i n these compounds except for the manganese(ll) and copper(ll) compounds. Cu(S0.jF)2 shows the 9 band spectrum t y p i c a l of C fluorosulfate groups, but i n Mn(S0_F) o the s p l i t t i n g of a l l the E modes are not resolved, and only 8 fundamental vibrations are observed. A further aim of t h i s work was to check the r e l i a b i l i t y of fluorosulfate frequency s h i f t s as measures of cation-anion interaction. We have shown i n t h i s work that while there i s a general increase i n frequencies f o r the Vg and v i b r a t i o n a l modes with'increasing cation-anion i n t e r a c t i o n , no quantitative relationship i s present. I f the vibra-t i o n a l spectrum i s to be used to determine the extent of t h i s i n t e r a c t i o n , then the band i s preferable to , since i t i s about twice as sensitive to changes i n the covalent in t e r a c t i o n . However, there are r e s t r i c t i o n s to i t s usefulness. The frequency does give a good in d i c a t i o n of the extent of cation-anion in t e r a c t i o n f o r compounds with predominantly i o n i c character. But where t h i s interaction i s more s i g n i f i c a n t , f o r example i n the bidentate or tridentate fluorosulfates, becomes rather insensitive to changes i n the degree of t h i s i n t e r a c t i o n , and i t correspondingly becomes less r e l i a b l e . While band frequencies may not always r e l i a b l y measure the degree of cation-anion in t e r a c t i o n i n a fluorosulfate compound, they may s t i l l be used together with the observed symmetry of the fluorosulfate group as an aid i n distinguishing between i o n i c , unidentate, bridging bidentate, tridentate, or quadridentate fluorosulfates. Compound Bonding Symmetry v 9(cm~ ) v f S p l i t t i n g (cm" ) Reference CsSO^F ionic 3 v 716 - 37 m(cSE^)H ( S O 3 F ) 2 semi-coordinated unidentate C s 71+1+ 75 t h i s work C ( S O 3 P ) ^ unidentate C s 81+3 21+2 70 F e ( S 0 3 F ) 3 "bridging bidentate C s 850 223 11+ N i ( S 0 3 F ) 2 bridging tridentate 3 v 859 - t h i s work C u ( S 0 3 F ) 2 bridging tridentate C s 861 83 ii it T i C 1 3 . 3 3 ( S ° 3 F ) 0 . 6 6 bridging quadridentate 3 v 650 - 71 Table VII - Bonding and Infrared Spectral Changes i n Fluorosulfates. 62 Table VII i l l u s t r a t e s some general differences i n the infrared spectra of fluorosulfate compounds with di f f e r e n t modes of bonding. Ionic fluorosulfate groups should possess symmetry except i n those compounds where the symmetry may be lowered by s i t e symmetry effects. The Vg band also occurs at low frequency, usually below 800 cm . As the covalent in t e r a c t i o n becomes stronger, there w i l l be a s h i f t to higher frequency f o r and the E modes may be s p l i t . In Ni(C^H^N)^(S0 3F)^, the trans fluoro-sulfate groups may be described as semi-coordinated unidentate ligands with la r g e l y i o n i c character. For the covalent C C S O ^ F ) ^ , the frequency i s considerably higher than f o r i o n i c fluorosulfates, and the s p l i t t i n g of the E modes i s much larger. Bidentate bridging fluorosulfate groups a l l show C symmetry, and they show large E mode s p l i t t t i n g s and s h i f t s to high s v*2 frequencies r e l a t i v e to i o n i c fluorosulfates. T y p i c a l l y tridentate fluorosulfate groups possess symmetry with the Vg frequency i n d i c a t i n g s i g n i f i c a n t covalent interaction. I f there i s departure from 0^ symmetry about the metal, then the fluorosulfate groups may show C symmetry, as we s observe i n copper(ll) fluorosulfate. In t h i s case, the d i s t o r t i o n of the CuOg octahedron i s a consequence of the Jahn-Teller effect. In the f i n a l p o s s i b i l i t y of quadridentate bridging fl u o r o s u l f a t e , we expect symmetry fo r f l u o r o s u l f a t e , and a d r a s t i c a l l y lowered frequency due to the bridging through the S-F bond. The only known example of t h i s type of fluorosulfate bonding i s i n T i C l ^ ^^(SO^F)^ gg, where the symmetry of the fluorosulfate group i s lowered s l i g h t l y due to s i t e symmetry effects. While we can often determine the bonding type of the fluorosulfate group simply from the symmetry i t shows i n the infrared spectrum, there may be some ambiguity i n distinguishing between fluorosulfates which have the same symmetry, but are bonded d i f f e r e n t l y . For example, Table VII shows 6 3 sim i l a r frequencies and s p l i t t i n g for C^O^F)^ and Fe^O^F)y though one contains unidentate fluorosulfate, and the other bridging bidentate fluorosulfate. Fortunately, the available l i t e r a t u r e on fluorosulfate compounds permits us to obtain the frequency ranges of the fluorosulfate stretching modes for the di f f e r e n t bonding types. These are presented i n Table V I I I , and represent the ranges i n which the large majority of the fluorosulfates i n each c l a s s i f i c a t i o n f a l l s . While the table shows the fluorosulfates neatly divided into d i s t i n c t bonding types, there i s not always such a clearcut d i s t i n c t i o n i n the compounds, and there w i l l always be some i n which the bonding i s intermediate between two types. With t h i s i n mind, the frequency ranges for the di f f e r e n t bonding types i n Table VIII w i l l be of some use i n assigning the v i b r a t i o n a l spectra of fluorosulfate derivatives, and also as a deductive aid i n determining the mode of bonding i n these compounds. The ranges given f o r unidentate fluorosulfate were obtained from the infrared spectra of some f o r t y - s i x compounds described i n the l i t e r a -27 28 70 72 ture . Those compounds whose frequencies do not f a l l within the ranges given a l l have low S-0 stretching frequencies. In some cases, the band positions may be lowered by intermolecular associations, f o r example, the hydrogen bonding i n f l u o r o s u l f u r i c acid. The halogen and interhalogen fluorosulfates often show low frequencies f o r V-(S0 o) and v (S0 o), and high frequencies f o r v ( S 0 * ) ^ n ) ' ^  '^3. For the bridging bidentate fluorosulfate ranges, we have studied the v i b r a t i o n a l spectra of thirteen compounds^'^' 2 2>2U»26,7li^ Q^ ^-^Q-^ have stretching frequencies within the ranges given i n Table V I I I . S i m i l a r l y , the frequency ranges f o r the Unidentate Bridging Bidentate C Symmetry Assignments Band Frequencies (cm ) Assignments Band Frequencies (cm ) » a < S 0 2 > 1^10 - 11+20 Va ( S ° 2 ) 11+20 - 1330 v s (so 2 ) 1280 - 1230 V S 0 2 .) 1250 - 1100 v(so ) 1010 - 790 v(so) 1090 - 1050 v(SF) 890 - 800 V ( S F ) 870 - 820 Symmetry Assignments Ionic Tridentate Quadridentate 1310 - 1250 1300 - 1250 1250 1110 - 1070 1120 - 1070 1070 V(SF) 810 - 700 860 - 820 6 5 0 Table VIII - Vibrational Frequency Ranges for Fluorosulfates 65 i o n i c fluorosulfates were obtained from the v i b r a t i o n a l spectra of twelve 37 75 compounds , and those of the tridentate fluorosulfates were determined from ten compounds. There i s only one known example of quadridentate fl u o r o s u l f a t e , and that i s T i C l - ..(SCLF)-3*33 3 '0.66 There are some general differences i n the band frequencies f o r unidentate and bridging bidentate fluorosulfate groups with C symmetry. s For unidentate fluorosulfate, the f i r s t two S-0 stretching modes occur at higher frequencies, and the t h i r d S-0 band occurs at lower frequencies than for bridging bidentate fluorosulfate. Since the nature of the cation and the other ligands coordinated to i t w i l l also influence the positions of the S-0 stretching modes, i t i s not surprising that a wide range of frequen-cies i s observed f o r the t h i r d S-0 band i n unidentate fluorosulfate group, which i s assigned to V(S-0 ), and for the f i r s t and second S-0 bands i n bridging bidentate fluorosulfate groups, which, according to Aubke and 2 l i 26 ^ coworkers are assigned to v (S0 Q ) and v (S0 o ) respectively. 0 represents the coordinated oxygen. For fluorosulfate groups with C^v symmetry, i t i s not possible to decide between i o n i c , tridentate and quadridentate fluorosulfates on the basis of the S-0 stretching frequencies. However, the mode of bonding can be determined from the S-F stretching frequency. The ranges f o r the bands are d i s t i n c t i n each case, with v(S-F) for tridentate fluorosulfate occurring at the highest frequencies, followed by i o n i c fluorosulfate, with quadridentate fluorosulfate occurring at very low frequencies because of the bridging S-F bond. The "frequency ranges presented i n Table VIII f o r the di f f e r e n t modes of fluorosulfate coordination may be useful i n assigning the v i b r a t i o n a l spectra of compounds where two types of fluorosulfates exist i n the same 66 molecule. Thus, t i n ( l V ) fluorosulfate contains both terminal unidentate 22 and bidentate bridging fluorosulfate groups, and with the frequency ranges i n Table VIII i t i s possible to assign the stretching modes. The assign-ments of the Raman frequency s h i f t s are given i n Table IX. -1 -1 Assignment Unidentate (cm ) Assignment Bidentate (cm ) vJSCv,) 11+31 V a ( S 0 2 * ) 1 1 + 2 0 V g(S0 2) 1233 V S ( S 0 2 * ) 1 1 2 i + v(S - 0 * ) 911 v( s_o) 1075 v(S - F ) 827 V ( S - F) 82+5 Table IX - Assignment of Stretching Modes for SnfSO^F)^. The s p l i t t i n g of the highest frequency E mode has been suggested 76 as an in d i c a t i o n of the type of the coordination occurring i n carbonato-, 77 and n i t r a t o - compounds. S i m i l a r l y , the separation between the symmetric and antisymmetric C-0 stretching frequencies has been used to dist i n g u i s h z?o 78 between unidentate and bidentate coordination i n haloacetates. ' When we consider the s p l i t t i n g of the highest frequency E mode i n the unidentate and bridging bidentate fluorosulfate groups, we f i n d that the magnitude of the s p l i t t i n g i n both cases i s comparable. For the unidentate fluorosulfates from which Table VIII was derived, there i s a 6 7 range of s p l i t t i n g from 253 cm to 159 cm" , whilst for bridging bidentate -1 -1 fluorosulfates, the range i s 395 cm to 100 cm . In view of the great s i m i l a r i t y of s p l i t t i n g s i n both cases, we must conclude that the magnitude of the s p l i t t i n g of the band i s not a good c r i t e r i o n for distinguishing between unidentate and bridging bidentate fluorosulfate groups. 68 CHAPTER h PYRIDINE COMPLEXES OF TRANSITION METAL FLUOROSULFATES  AND RELATED COMPLEXES U.1 INTRODUCTION As was mentioned i n Chapter 1, much of the work on coordination complexes of fluorosulfate s a l t s \wa'S2 done when modern techniques of characterization were not available. Even more recently, there has been very l i t t l e systematic work to determine the nature and strength of the fluorosulfate coordination. 29 Milne prepared the t e t r a k i s ( a c e t o n i t r i l e ) complexes of copper(l), c o p p e r ( l l ) , and z i n c ( l l ) fluorosulfate whose infrared spectra indicate a strong i n t e r a c t i o n between the fluorosulfate anion and the complex copper(ll) and z i n c ( l l ) cations. He concluded from e l e c t r i c a l conductivity measurements i n a c e t o n i t r i l e that there was strong association between the fluorosulfate anion and the d i p o s i t i v e cation, but the copper(l) complex shows no strong i n t e r a c t i o n with fluorosulfate anions. 30 Mayfield has also prepared a series of complexes of the formula JcuL^SO^F^J (L = pyridine, l+-methylpyridine, U-aminopyridine), but did not succeed i n preparing si m i l a r complexes with other t r a n s i t i o n metal fluorosulfates. His studies, which are the most complete so f a r reported, 69 indicate coordination of the fluorosulfate group to form tetragonally distorted octahedral complexes. He suggests from infrared studies of the copper-nitrogen vibrations i n the tetrakis(pyridine) complexes, that flu o r o s u l f a t e , perchlorate, perrhenate, and hexafluorophosphate ions are very s i m i l a r i n t h e i r bonding behaviour, but he makes no attempt at deter-mining the coordination strength of the fluorosulfate ion r e l a t i v e to the other anions. He observes from powder x-ray d i f f r a c t i o n data, that Cu(py)^(S0 3F)2 i s i s o s t r u c t u r a l with the corresponding perchlorate complex, but not with the perrhenate complex. In view of the scar c i t y of information on the bonding character-i s t i c s of the fluorosulfate ion, i t seemed desirable to study a series of complexes of t r a n s i t i o n metal fluorosulfates and compare them with other complexes which d i f f e r only i n the anion involved. Since quite extensive investigations on the coordination of univalent unidentate anions i n pyridine complexes have been reported pre-viously, 3 1''' 3^'^'' ' 7^~^ 3 i t was decided to prepare pyridine complexes of t r a n s i t i o n metal fluorosulfates f o r easy comparison. Although much i s known about pyridine complexes of t r a n s i t i o n metal s a l t s such as perchlorates, tetrafluoroborates, n i t r a t e s , and t r i f l u o r o -acetates, very often the published information has been reported by various workers who studied the complexes under di f f e r e n t conditions using different techniques. Comparison of a large body of data from many sources opens up the p o s s i b i l i t y of r e s u l t s which may vary simply due to differences i n experimental c a l i b r a t i o n s , or systematic errors. In. order to minimize these sources of error, much of the previous work on selected complexes has been repeated so that any systematic error would be the same for a l l complexes. 70 As the complexes are a l l studied under the same conditions, we can now draw more r e l i a b l e conclusions from differences i n the physical properties of the complexes. To f a c i l i t a t e comparison of the coordination of the fluorosulfate ion with the other anions, the copper(ll) tetrakis(pyridine) complexes were selected f o r detailed study. This series of complexes were chosen for Q several reasons: ( i ) pure compounds are e a s i l y prepared, ( i i ) the d elec-tron configuration of copper(ll) allows for easy interpretation of electronic and electron paramagnetic resonance spectra, and magnetic s u s c e p t i b i l i t y data, ( i i i ) much information i s available on Cu(py)^X 2 complexes. An additional value of re-examining previously studied t e t r a k i s (pyridine) complexes i s that, while magnetic s u s c e p t i b i l i t y , infrared and electronic spectra have usually been reported, we have the opportunity of repeating these studies under a greater range of conditions. We can also include less commonly used techniques such as electron paramagnetic reson-ance spectroscopy and e l e c t r i c a l conductivity which can provide useful information on the complexes themselves. This work represents the f i r s t systematic and extensive study of complexes of the type Jcu(py)^X 2J (X = SO^F-, BF^~, C10^ , NO^ ", CH^CgH^SO^", CF^COCT). We have studied these complexes by i n f r a r e d , electronic, and electron paramagnetic resonance spectroscopy both i n the s o l i d state and i n solution. We have also measured e l e c t r i c a l conductances i n a v a r i e t y of s o l -vents and the magnetic s u s c e p t i b i l i t i e s over a temperature range. Many of these r e s u l t s are reported here for the f i r s t time. In p a r t i c u l a r t h i s i s the f i r s t reported preparation and characterization of Ni(py)^(S0 3F) 2, Zn(py)^(S0 3F) 2, and C u t p y ^ C H ^ H ^ S O ^ . 71 it.2 INFRARED SPECTROSCOPY 1+.2.1 Anion Bands i n Fluorosulfate Complexes The fluorosulfate ion possesses C^v point symmetry with s i x fundamental vibrations which are infrared and Raman active. When i t i s coordinated to a metal ion, there may be a reduction i n symmetry to C s (Figure 1+) and a s h i f t i n band positions (Table V i i ) r e l a t i v e to those observed i n the spectra of io n i c s a l t s such as j^^AsJsO^F. The infrared band assignments for the fluorosulfate group i n n i c k e l ( l l ) , c o p p e r ( l l ) , and z i n c ( l l ) tetrakis(pyridine) fluorosulfates are l i s t e d i n Table X, and are given i n terms of C^V symmetry. The appearance of the infrared spectra of the complexes are t y p i f i e d by Ni(py)^(S0 3F)2 which i s shown i n Figure 13» The pyridine complexes studied show that the symmetry of the fluorosulfate group i s lowered to C with a l l observable E modes s p l i t . s The bands a r i s i n g from the E mode are not observed i n the infrared spectra since they are expected to have f a i r l y weak i n t e n s i t i e s , and i n a l l p r o b a b i l i t y , are hidden by the 16b out-of-plane pyridine deformation band which occurs i n the same region. The infrared spectrum of Cu(py)^(S0 3F)2 30 obtained i n t h i s work agrees well with that reported previously, though there i s a difference i n the assignment of the stretching vibrations. The s h i f t i n the band i n these complexes to higher frequencies r e l a t i v e to " i o n i c " f l u o r o s u l f a t e , e.g. ^ ^AsJsO^F, together with a reduction i n symmetry of the fluorosulfate groups suggests that the SO^F ions are involved i n coordination to the metal ions as unidentate ligands. The magnitude of the E mode s p l i t t i n g and the s h i f t i n to higher 73 ^ ( E ) "-| ( A ) ^ 2 ( A ) ^ ( E ) ^ ( A ) Ni(py) i +(S0 3F) 2 1333s 1259s 1 0 8 2 s 7 4 3 s 5 9 0 s 5 7 3 s 557m Zn(py) i +(S0 3P) 2 1331s 1259s 1 0 8 0 s 7 3 7 s 5 8 8 s 5 7 4 s 557m Cu(py)^(S0 3P) 2 1316s 1 2 6 6 s 1072s ^ 7 3 2 s 5 8 0 s s h 5 7 8 s 558m _ ^ 7 2 4 s _ _ ^ A s S O ^ 3 7 1288 1068 7 0 4 5 7 4 5 7 0 s h C E ^ S O ^ 1465 1235 787 840 Table X - Infrared Band Assignments (cm ) for M(py)^(S0 3F) 2 i n the Sol i d State. frequency intimates a "semi-coordination" of the fluorosulfate groups. 85 86 This term has been used to describe anions which are involved i n coordination and of s u f f i c i e n t strength to produce s l i g h t d i s t o r t i o n s of the anion, but which are not s u f f i c i e n t l y strong to produce the gross d i s t o r t i o n and f u l l symmetry lowering which occurs when the anions are involved as f u l l y coordinated groups. Table X shows-the s p l i t t i n g of the band i n the complexes to be 50 - 74 cm . For a compound where the fluorosulfate group i s indes-putedly bonded i n a f u l l y coordinated unidentate fashion, e.g. methyl -1 f l u o r o s u l f a t e , CILSO-F, there i s a much greater s p l i t t i n g of 230 em . ^ * could While the s p l i t t i n g of the E modes i n the complexes I conceivably be due to s i t e symmetry eff e c t s , e l e c t r i c a l conductivity measurements i n dichloro-methane solution (section 4 - 3 ) also show that the S0 3F groups are coor-dinated and there i s very l i t t l e anion di s s o c i a t i o n i n t h i s solution. C H 2 C , 2 1600 1400 1200 1000 Wavenumber (cm ) 8 0 0 6 0 0 4 0 0 Figure Mi - Infrared Spectrum of Cu(py)^(S0 3F )o i n Dichloromethane Solution. 75 The only changes i n the infrared spectra of Ni(py)^(S0 3F) 2 and Cu(py)^(S0 3F) 2 i n dichloromethane solution from the s o l i d state spectra are a very small s h i f t i n the stretching modes to lower frequencies, and a removal of the s p l i t t i n g of some pyridine hands. This suggests that i n dichloromethane, which i s a poorly coordinating solvent, there i s very l i t t l e or no d i s s o c i a t i o n of the fluorosulfate groups. No di s s o c i a t i o n of pyridine could be detected, even i n concentrated solutions. Figure 1L\ shows the infrared spectrum of Cu(py)^(S0 3F) 2 i n dichloromethane solution. The addition of up to 2$% pyridine to the solution causes very l i t t l e change i n the infrared spectrum of Ch^py^SO^F^. Ni(py)^(SC> 3F) 2 seems to show a decrease i n the magnitude of s p l i t t i n g i n the E modes as the proportion of pyridine i n the solution i s increased.. This i s not completely unambiguous from the bands since the lower frequency compo-nent of the s p l i t i s p a r t i a l l y obscured by a strong absorption of the dichloromethane solvent centered at 1265 cm . This solvent band also makes i t impossible to detect the presence of ionized fluorosulfate groups because i t would completely overwhelm any small fluorosulfate absorption by v i r t u e of i t s i n t e n s i t y and frequency. Ni(py)^(S0 3F) 2 and Cu(py)^(S0 3F) 2 are sparingly soluble i n pyridine and the infrared spectra of the saturated solutions show fluoro-sulfate groups with symmetry. The only c l e a r l y resolved bands i n the spectra are those of the vibrations which occur as sharp bands at 1283 —1 —1 cm" and 1286 cm" for the n i c k e l ( l l ) and copper(ll) complex respectively. This suggests that i n pyridine solution, the species are JM(py)gj 2 + and i o n i c SO^F". The hexakis(pyridine) n i c k e l ( l l ) cation has been previously observed for n i c k e l ( l l ) tetrakis(pyridine) perchlorate and tetrafluoro-r , + 3 5 , 8 0 , 8 1 borate. 76 In . acetonitrile solvent, Ni(py)^(S0 3F) 2 and Cu(py)^(SO.jF) appear to dissociate anions completely to give free fluorosulfate ions possessing symmetry. The infrared bands of the fluorosulfate group are tabulated i n Table XI and are compared with ionic j^^AsJsO^F i n the solid state. vk V1 v2 v$ v 3 v 6 Ni(py ) 1 + ( S 0 3F) 2 1281; 1066 701; Slk Cu( Py ) ^ ( S 0 3F) 2 1282 1067 707 577 [ V s ] S0 3F 1288 1068 701; 571; 5 7 0 s h Table XI - Band Assignments (cm" ) for Ni(py)^(S0 3F) 2 and  Cu(py)^(S0 3F) 2 i n Acetonitrile Solution. Evidence for or against pyridine dissociation could not be detected because of intense solvent absorption. When Ni(py)^(S0 3F) 2 i s dissolved i n nitromethane, there appears to be a slow dissociation of fluorosulfate ions. A fresh solution of the complex shows bands at 1261 cm (the other component of the s p l i t i s —1 _ hidden by solvent absorptions), 1282 cm (due to ionized S03F~ with C 3 v symmetry), and 760 cm . I f the solution i s allowed to stand for several -1 -1 days, there i s a decrease in intensity of the bands at 1261 cm and 7^0 cm _1 concomitant with an increase i n intensity of the 1282-cm, band and the 77 _1 appearance of a strong new band at 737 cm . Some dis s o c i a t i o n of pyridine i s also evident from the infrared spectrum of the solution. Cu(py)^(S0 3P)2 shows si m i l a r behaviour i n nitromethane solution, and there appears to be more extensive pyridine d i s s o c i a t i o n than i n the ni c k e l complex. 80 Rosenthal and Drago have also observed anion displacement i n nitromethane solutions of the n i c k e l ( l l ) tetrakis(pyridine) perchlorate and tetrafluoroborate. U.2.2 Anion Bands i n Other Complexes  Perchlorate The free perchlorate ion belongs to the symmetry point group T^ . The f i v e atoms contribute nine degrees of freedom, but due to the high symmetry of the ion, there are only four normal modes. Of these modes, only (asymmetric stretching), and (asymmetric bending) are infrared active. Upon coordination, the T^ symmetry beci=sie reduced to symmetry which r e s u l t s i n a s p l i t t i n g of the two t r i p l y degenerate normal modes of vi b r a t i o n and an a c t i v a t i o n of the two symmetric modes. The correlation table f o r T^ and 0^ point groups i s shown i n Table XII. The band assignments f o r the perchlorate complexes prepared i n t h i s work are given i n Table X I I I . The infrared spectra of the perchlorate complexes are i n good agreement with those reported previously. They show that the perchlorate groups are acting as unidentate ligands of symmetry C^. Zn(py)^(C10^)2 alone shows T^ symmetry f o r perchlorate, with the metal also i n tetrahedral 3) coordination. This i s i n contrast with Zn(py)^(S0^)^ whose infrared 78 State of Anion CIO. Symmetry T, -0-C10- C_ 3 3 v K 1 A(R) "2 E(R) sym.str. sym.bend y2 A^I.R) C10*str. " 3 F 2(I,R) asym.str. "6 "1 E(I,R) A(I,R) rock sym.str. CIO^ asym.str. CIO^ F 2(I,H) asym.bend V 3 A(I,R) E(I,R) sym.bend asym.bend CIO, CIO, Table XII - Vibrations of the Perchlorate Group as a Function of Symmetry. spectrum, and e l e c t r i c a l conductivity i n dichloromethane solution (section i+.3) indicate an octahedral stereochemistry f o r the zinc ion with coordi-nated fluorosulfate groups. The difference i n anion coordination i n the zinc complexes therefore suggest that the fluorosulfate group i s a stronger ligand than the perchlorate group. The magnitude of the s p l i t t i n g i n the Fg1modes, and the Vg frequency i n the complexes compared with HCIO^ which contains unidentate perchlorate groups, suggests a weak coordination of CIO^ . S p l i t t i n g s i n _-| the band (T^ symmetry) of 101;, 1 0 1 , and 58 cm" f o r the c o b a l t ( l l ) , n i c k e l ( l l ) , and copper(ll) complexes are small compared with the s p l i t t i n g 79 v 2(A) V^(E) v 3(A) v 6 ( E ) [ c o ( p y ) i + ( C 1 0 i + ) J 1 1 3 7 s 103l+s 9 3 1 m 6 1 7 m -[ u i ( p y ) u ( C 1 0 u ) 2 ] 1 1 3 3 s 1 0 3 2 s 9 3 1 m 6114m -[ c u ( p y ) ^ ( C 1 0 u ) J 1 1 1 3 s 1 0 5 5 s 9 3 3 m 6 2 5 m 6 2 0 m -1 1 0 0 - - 1 0 6 0 91+5 6 2 2 -8 7 1 1 1 0 9 3 2 6 2 6 1+60 HCIO^ 8 7 1 3 1 2 1 0 3 2 7 3 9 5 8 5 571+ 1+26 pyridine band overlap in this region. Table XI I I - Band Assignments (cm ) f o r M(py)^(C10^) 2 i n the Solid State of 2 8 0 cm i n HCIO^. In addition, the frequencies are closer to i o n i c values than covalent. The infrared spectra are therefore consistent with a semi-coordination of the perchlorate groups i n the tetrakis(pyridine) complexes studied. The infrared spectrum of Cu(py)^(C10^) 2 i n dichloromethane solution shows no change i n the perchlorate bands from the spectrum of the complex i n the s o l i d state. There i s no evidence f o r dis s o c i a t i o n of perchlorate groups or pyridine i n t h i s solvent. 80 T e trafluoroborate The tetrafluoroborate ion has the same symmetry as the perchlorate ion and also lowers i t s symmetry from T^ to upon coordination as a unidentate ligand. The band assignments f o r Cu(py)^(BF^) 2 given i n Table XXV i n terms of symmetry. v^(E) v., (A) v 2(A) v^(E) v 3(A) v g(E) Cu(py)^BP^) 2 1073sb 9 8 5 s 761 s 518m 88 B P " 98U 769 521+ 353 — 1 Table XIV - Band Assignments (cm" ) for Cu(py)^(BF^) 2 i n the S o l i d State. As was observed f o r the perchlorate complexes, the infrared spectra of Cu(py)^(BF^) 2 showu the tetrafluoroborate groups to be semi-coordinated to the metal ion. There i s a small s p l i t t i n g of the v ^ F ^ band i n T^ symmetry, and the frequencies of the infrared bands are simil a r to those of io n i c tetrafluoroborate. 81 The infrared spectrum of Cu(py)^(BF^) 2 has been reported by Brown et a l . , and the spectrum obtained here agrees well with t h e i r r e s u l t s . -1 -1 They assign bands at 1+37 cm to Vy and 517 cm to Vy This implies that upon coordination, the band i n T^ symmetry s p l i t s by 82 cm" (!) while the stretching mode s p l i t s by only 95 cm . I t i s reasonable to expect that the r e l a t i v e magnitudes of the s p l i t t i n g of tetrafluoroborate E modes should be comparable with those of perchlorates or fluorosulfates with si m i l a r bonding. Cu(py)^(C10^) 2 shows a s p l i t t i n g i n the v ^ ( F 2 , T^) mode of 58 cm while the s p l i t t i n g i n " l ^ ( F 2 , T^) i s only 5 cm" . In Cu(py)^(S0 3F) 2, the s p l i t t i n g of v^(E, C^) i s too small to be completely resolved. Since the degenerate stretching mode i n the perchlorate and fluorosulfate complexes shows such a large s p l i t t i n g r e l a t i v e to the degenerate bending mode, i t 3 82 3k _1 appears u n l i k e l y that there should be a s p l i t t i n g as large as 82 cm f o r the v^(^2' ^d) m o < i e i n Cu(py)^(BF^) 2. The band at 2+37 cm i s more reasonably assigned to the pyridine l6b out-of-plane bending mode. Brown 3ii —1 et. a l . report a band at 1+62 cm i n the spectrum of Cu(py)^(BF^) 2 which they suggest could be associated with the copper-fluorine bands. No such band was observed i n t h i s work. p-Tolyl4u^a%e-— -The point symmetry about the sulfur atom i n the p - t o l y l s u l f a t e (Ts) group, considering only those atoms d i r e c t l y bonded to i t , i . e . , the -CSO^ group, may be taken to be for which we expect s i x infrared and Raman active vibrations. When the p - t o l y l s u l f a t e ion i s coordinated as a unidentate ligand, the symmetry i s lowered to C and the doubly degenerate E modes are s s p l i t to give nine infrared and Raman active bands. 82 The band assignments for the p - t o l y l s u l f a t e group i n the C u ( p y ) ^ ( l l s ) 2 complex are l i s t e d i n Table XV. The assignments are given i n terms of symmetry, and the descriptions are s i m i l a r to those f o r the fluorosulfate group (Figure 3 ) ' i / ^ 4 ( E ) V A ) V 2 ( A ) r ' ^ E ) V 3(A) V g(E) Cu(py) i |(Ts) 2 1 2 4 0 s 1185s 1 0 3 5 s 6 8 1 s 5 6 8 s 5 6 5 s 549m 89 Ts" 1125 1030 c i y r s 9 0 1352 1188 1177 Table XV - Band Assignments (cm~ ) f o r Cu(py)^(Ts) 2 i n the S o l i d State. Very l i t t l e work has been done on the infrared spectra of 89-92 p - t o l y l s u l f a t e s , and the few published reports seem content to simply l i s t the bands without assignments, or to assign only the highest frequency bands. The p - t o l y l s u l f a t e group was assigned here by comparison with the infrared spectra of the fluorosulfate and methylsulfate groups. As we observe for the other anions, the infrared spectrum of Cu(py)j^(Ts) 2 shows a reduction i n symmetry consistent with unidentate p - t o l y l s u l f a t e groups. The s i m i l a r i t y of band frequencies with i o n i c 89 p - t o l y l s u l f a t e and the small s p l i t t i n g of E modes indicate weakly coordinated p - t o l y l s u l f a t e groups. 8 3 Dichloromethane solutions of Cu(py)^(Ts)2 show great propensity for anion exchange with KBr c e l l windows. For good spectra, i t i s necessary to p o l i s h the c e l l surfaces between each run, and to minimise the time between introduction of the solution into the c e l l and the actual recording of the spectrum. I f these precautions are carried out, then the infrared spectrum shows s l i g h t d i s s o c i a t i o n of pyridine from the complex, but there i s l i t t l e change i n the band positions of the p - t o l y l s u l f a t e group. Addition of 10% pyridine suppresses the pyridine d i s s o c i a t i o n with no evidence of p - t o l y l s u l f a t e d i s s o c i a t i o n . The infrared spectra of pyridine solutions of Cu(py)^(Ts)2 s t i l l show p - t o l y l s u l f a t e groups with C symmetry though the presence of a strong s pyridine band i n the region of the asymmetric SO^ stretching modes d i s t o r t s the band shapes and i n t e n s i t i e s of the stretching vibrations. In spite of _-| t h i s , there i s some ind i c a t i o n of a new band appearing about 122J? cm which may arise from some s l i g h t anion dis s o c i a t i o n from the complex to give p - t o l y l s u l f a t e ions with Q symmetry. Nitrate The n i t r a t e ion of D^ ^ symmetry gives r i s e to four fundamental vibrations, only one of which i s infrared forbidden. On destroying the threefold symmetry by coordination, the symmetry of the group i s lowered to degeneracy of the E modes i s l i f t e d . A l l modes become infrared active to give s i x bands i n the infrared spectrum. The vibrations of the n i t r a t e group as a function of symmetry are shown i n Table XVI. 8k State of Anion NO. Symmetry -0-N0„ "1 A^R) sym. s t r . A 2 ( I ) o.p.def. ; 2 v " 2 A^I.R) NO*str. E (I,R) asym.str. v6 v 1 vi+ B 2(I,R) A^I.R) B^I.H) rock sym.str. asym.str. V k E-'(I,R) i.p.def. v 3 " 5 A^I.R) B.,(l,R) sym.def. asym.def. NO, N0r Table XVI - Vibrations of the Nitrate Group as a Function of Symmetry. v ^ ) v ^ ) v 2 ( A l ) v 6 ( B 2 ) v 3 ( A 1 ) V ^ ) Cu(py)^(N0 3) 2 1l+07s 1 3 0 9 s 10U6s 823m 7 0 9 * N0 3" ^ "l390 ' 1050 831 720 CH"3N03 9 ^ 1 6 7 2 V S 1287s CQBke 759m * Choca, Ferraro, and Nakamoto calculate = 730 cm" , = 718 cm from combination bands. Table XVII - Band Assignments (cm ) f o r Cu(py)^(N0 3) 2 i n the S o l i d State. 85 The band assignments of the n i t r a t e groups i n Cu(py)^(N0 3) 2 are given i n Table XVII i n terms of C 2 v symmetry and the band positions are compared with those of i o n i c and covalently bonded unidentate groups. The infrared spectrum of Cu(py)^(N0 3) 2 i n the s o l i d state has been previously studied and good agreement i s obtained here with the e a r l i e r s t u d i e s . 9 3 ' ^ 1 ' 1 The s p l i t t i n g i n the band (E , D^) of 98 cm" f o r —1 Cu(py)^(N0 3) 2 > compared with 3 8 5 cm i n covalently bonded n i t r a t e groups (CH^NO^) indicates only a small d i s t o r t i o n from symmetry for n i t r a t e i n the complex. Combined with t h i s , the s i m i l a r i t y of the band frequencies with those of i o n i c n i t r a t e suggests that the n i t r a t e group i s weakly coordinated i n Cu(py)^(M) 3) 2. The infrared spectra of dichloromethane solutions of Cu(py)^(lT0 3) 2 show extensive d i s s o c i a t i o n of pyridine with the formation of the b i s p y r i -dine complex, Cu(py) 2(N0 3) 2. There i s a considerable decrease i n i n t e n s i t y -1 -1 of the band at 11+07 cm and strong new bands appear at 1474 cm and -1 1279 cm . Such a change i s i n agreement with the spectrum reported by 93 -1 -1 Ferraro et a l . who observed N-0 stretching bands at 11+74 cm and 1282 cm f o r Cu(py^ 2(N0 3) 2. _-| On addition of 10% pyridine to the solution, the bands at 11+74 cm and 1279 cm disappeared and the N-0 stretching region shows a closer resemblance to that of the tetrakis(pyridine) complex although a new band appears at 1355 cm whose o r i g i n i s uncertain. For Cu(py) 3(N0 3) 2, the only -1 -1 93 n i t r a t e bands expected i n this, region are at 1458 cm and 1291 cm . 86 Trifluoroacetate The v i b r a t i o n a l spectrum of the trifluoroacetate ion i s consis-i tent with C point symmetry, f o r which we expect nine frequencies i n the A s IT class and s i x i n the A class. The observed infrared frequencies and assignments f o r the trifluoroacetate group i n Cu(py)^(CF 3C00)2 are l i s t e d i n Table XVIII. The band assignments used are those of Crowder and Jackson.^ 6 97 96 Assignment Cu(py)^(CF 3C00) 2 CF^OO" CF^COOCB^ 10 11J v. 12 C = 0 s t r . 0 s t r . CF^ s t r . C-C s t r . 0C0 def. CF^ def. (o.p.) (i.p.) (o.p.) C-C09 def. (o.p.) 1689 s 1681 1793 1670 ssh 11+16 m 11+35 1207 1205 s 1202 1229 1177 s 1178 1116 s 111+3 111+5 828 s 81+1+ 81+2 797 s 729 780 ( 720 s 731+ | 595 vw 601 586 \ 516 vw 520 395 w 1+10 1+22 -1 Table XVIII - Band Assignments (cm ) for Cu(py)^(CF 3COO) 2 i n the S o l i d State 87 The infrared spectrum of Cu(py)^(CF 3COO) 2 has been previously 82 9 8 reported i n part by Agambar and O r r e l l , and Kettle and P i o l i . The band frequencies show good agreement with the e a r l i e r work, but we disagree 82 -1 with Agambar and O r r e l l ' s assignment of a band at 11+50 cm to the C-0 stretching mode. In a l l l i k e l i h o o d t h i s band i s due to the 19b r i n g v i b r a t i o n of pyridine. In dichloromethane solution, the infrared spectrum shows that Cu(py)^(CF 3C00) 2 dissociates pyridine extensively to give the Cu(py) 2(CF 3C00) 2 complex with bidentate trifluoroacetate groups. For bidentate CF^COO" with C symmetry, three C-F stretching frequencies, v„, V 0, V..., are expected, with and degenerate. For unidentate coordination, t h i s degeneracy i s removed. The infrared spectrum of Cu(py)^(CF3C00),2 i n dichloromethane shows only two C-F stretching bands at 1198 and 111+5 cm . Agambar and 82 O r r e l l have found that the band assigned to the C-C stretching v i b r a t i o n occurs at higher frequency i n the bis(pyridine) complex than i n the t e t r a k i s (pyridine) complex. The solution spectrum shows t h i s band at 81+3 cm" , compared with 828 cm" f o r Cutpy^CF^COO^ i n the s o l i d state. When pyridine i s added to the dichloromethane solution, the i n t e n s i t y -1 -1 of the 81+3 cm band decreases and a new band appears at 826 cm , which i s consistent with the conversion of some of the bis(pyridine) complex back to Cu(py)j^(CF 3C00) 2. In solutions containing 10% pyridine, the infrared spectrum s t i l l shows only two C-F stretching modes, in d i c a t i n g that there i s s t i l l l a r g e l y Cu(py) 2(CF 3C00) 2 i n solution. I t i s i n t e r e s t i n g that while the spectrum of Cu(py) 2(CF 3C00) 2 i n the s o l i d state shows two bands at 1696 and 1666 cm , the solution spectrum -1 shows only a single band at 1680 cm . Such a s p l i t t i n g of the C = 0 88 82 stretching band has been suggested to be a r e s u l t of v i b r a t i o n a l coupling between the two trifluoroacetate groups coordinated to the copper ion. Cu(py)^(CF 3C00)2 i s s l i g h t l y soluble i n pyridine, but the infrared spectrum of the saturated solution i s concentrated enough to show that the trifluoroacetate groups are s t i l l coordinated i n t h i s s o l ution; Three CF^ _-] stretching bands are observed at 1 2 0 2 , 117i+, and 1125 cm . The C-C stretching band also appears at 823 cm" . Thus Cu(py)^(CF 3C00) 2 shows no evidence of anion d i s s o c i a t i o n i n pyridine solution. I4..2.3 Pyridine Bands 99 G i l l and coworkers found that there i s a one-to-one correspondence i n the v i b r a t i o n a l spectra between complexed and free pyridine, and that each band i n the spectrum i s f a i t h f u l l y reproduced with only minor s h i f t s or s p l i t t i n g s . The assignment of the various bands can therefore be made by a straightforward comparison between the spectrum of coordinated pyridine and the spectrum of the free base, whilst most of the v i b r a t i o n a l bands are l i t t l e changed on coordination, the positions of the bands at 1578» 6 0 1 , and I4.O3 cm i n free pyridine are sensitive to the p o l a r i s i n g power of the metal and are shifted about 30 + 5 cm to higher frequencies i n the f i r s t row t r a n s i t i o n metal complexes. A l l the bands observed f o r the coordinated pyridine i n the complexes studied i n the 1650 - i+00 cm region are l i s t e d i n Table XIX. The band 99 assignments are based on those of G i l l and coworkers. Of the complexes l i s t e d i n Table XIX, only the spectra of the copper(ll) complexes of the 30 99 82 f l u o r o s u l f a t e , tetrafluoroborate, and trifluoroacetate have been reported previously i n any d e t a i l . There i s good agreement with these r e s u l t s , though more s p l i t t i n g of bands are observed here. 1 + 6b or 6a + 12 8a 8b 19a 19b 9a 15 18a 12 1 5 1) 11 6b 6a 16b K i ( p y ) [ t ( S 0 3 F ) 2 16IJ2 w 1609 a 1605 8 1572 vw 11)89 m 11)1)7 s 1237 m 1222 s 1156 m 1072 a 101)1) 8 1017 m 1011* m 955 w 833 v 761 B 701 a 651 vw 631 m 1*32 m Cu(py) | t(S0 3P) 2 161+1 vw 1610 s 1608 s 1572 vw 11)90 m 11)53 s 114,9 s 1239 m 1222 m 1160 m 1072 a 101)7 m 1020 m 1017 m 955 vw 876 w 761 S 696 s 650 vw 6!)3 m 639 m 1*33 m 1)28 w Zn(py), <(S0 3P) 2 1637 w 1609 8 1601) 8 1573 vw 11+90 m 11)51 ash 11)1)8 s 1238 m 1222 s 1159 m 1072 a 101)3 a 1018 m 1013 m - 760 S 700 s 651 vw 631) m 629 ra 1)22 m Co(py) l 4(C10 i <) 2 161*0 vw l607ssh 1601) s 1571* v 11)90 m 11)50 ash 114)8 8 1237 w 1221) s 1163 s 1072 8 - - 951) vw 760 8 702 a - 631 m 1)25 0 i)l8 wsh Ni(py) l 4(C10 [ i) 2 161)0 vw 1609 a 1596 s 1573 w 11)91 m 11)51 ash 114)8 a 1238 w 1225 8 1161) 8 1072 s - -951+ vw 881) vw 760 B 702 a 651 vw 632 m 1*33 m Cu(py)lt(ciolt)2 -1611 m 1609 m 1573 vw 11)92 m 11+5U a 114)9 a 121*0 w 121*5 a 1165 m 1072 a 101)6 s 1018 m 957 w 882 vw 762 8 759 a 696 s 650 w 61)5 w 61)0 m 1)36 m 1*29 w Cu(py) u(BP 1 ;) 2 -1612 s 1609 s 1573 vw 11)92 m 11)55 a 11)50 a 1238 w 1221+ 8 1161) m 1072 s 101)6 8 1017 s 881) vw 761 8 696 s 651 w 614) m 61)0 m 1)36 m 1)28 m Cu(py) l t(N0 3) 2 -1609 m 1602 m -11)97 m 11)86 m 11)53 a 114)2 s 1238 m 1227 m 1212 m 1163 m 111)7 m 1078 m 1068 m 1065 m 101)6 s 1017 m 970 w 893 w 782 m 769 m 762 s 701 s 651 vw 61)1 s 1*1*1* m 1)36 m Cu(py), <(OTs) 2 - 1607 s 1573 vw 11)90 m 11)52 a 1238 w 1225 8 1165 a 1155 a 1071) m 101)7 m -960 w 896 w 769 8 762 m 703 8 651 w 61)2msh 639 m 1)36 m 1*27 w Cu(py) u(CP 3COO) 2 - 1607 8 1576 w 1li9l) m 11)88 m 11*51 e 121)3 m 1227 a 1163 Q 1079 m 1073 m 1070 m 101)5 m 1018 m 991 w 960 w 766 n 762 8 759 B 700 a 651 vw 639 m 1*37 m 1*28 w TABLE XTX - TTnTTiTWP, PANT) ASPTHNTWrS TN M f n V X, , COMPLEXES 90 It has been claimed''that the shifts observed for the 6a and 16b pyridine ring deformation modes i n the complexes with metal halides are independent of the halogen and follow the Irving-Williams order for a series of complexes of the same stereochemistry differing only i n the central metal. The same order was found to exist i n the tetrakis(pyridine) complexes of the f i r s t row divalent transition metal perchlorates and tetrafluoroborate, 30 101 perrhenates, and thiocyanates. For the series of transition metal tetrakis(pyridine) complexes studied here, no similar trend can be found, since the 6 a and 16b pyridine bands indicate different orders of complex s t a b i l i t i e s . Anion, X CF-C00 CH0CvH. SO-,- NO ~ S0oF~ HP," CIO," 6 a ring deformation 639 61+2 sh 61+1 6I+3 61+1+ 61+5 639 639 61+0 61+0 Table XX - Pyridine in-plane Ring Deformation ( 6 a ) for Cu(py)^X2 Complexes I t i s also of interest to see whether the 6 a and 16b pyridine vibrations i n the complexes are independent of the a x i a l ligands i f the central metal ion i s kept constant. As the a x i a l ligand f i e l d strength weakens, the in-plane strength ought to increase due to increasing involve-ment of the copper o r b i t a l s with the in-plane ligands. Thus, i f there i s a dependence of the pyridine bands on the a x i a l ligands, then we would 91 expect a s h i f t to higher frequency i n the infrared spectrum for the most weakly coordinating anions. For the 6 a mode, we do observe a small v a r i a t i o n , as Table XX shows. I t therefore appears that the 6 a pyridine band i n complexes with weakly coordinating ligands such as perchlorate or tetrafluoroborate does occur at s l i g h t l y higher frequencies than i n those complexes with stronger ligands such as t r i f l u o r o a c e t a t e . However, i n view of the small differences i n frequencies involved, l i t t l e use can be made of the 6 a band i n determining the r e l a t i v e coordinating strengths of d i f f e r e n t anions. The infrared spectra of the tetrakis(pyridine) complexes studied show well defined s p l i t t i n g of some of the bands. These s p l i t t i n g s are often observed i n pyridine complexes, and are most l i k e l y due to interaction between molecules i n the unit c e l l , low l a t t i c e s i t e symmetry for the complex, or s l i g h t rotations of the coordinated pyridine about the metal-nitrogen bond. As we have mentioned e a r l i e r , the bands at 1 5 7 8 , 601 and I4.O3 cm i n free pyridine show the greatest s h i f t on coordination to a metal ion. Measurement,r,of the frequencies of these bands can therefore r e a d i l y d i s t i n -guish between coordinated pyridine and the free base. The presence of a -1 -1 band at 1067 cm i n addition to those at 1 5 7 8 , 601 and i|03 cm have been commonly used i n determining whether free base i s present together with the 7 6 coordinated ligand. I f there i s s i g n i f i c a n t amount of uncoordinated pyridine present, i t may even be possible to see a doubling of the t r i p l e t of bands ( 9 a , 15» 18a) between 1230 and 1070 cm i n addition to the more widely used bands at 1 5 7 8 , 601 and I4.O3 cm . 92 U . 3 ELECTRICAL CONDUCTIVITY Before t h i s work was i n i t i a t e d , only one previous investigation was made of the e l e c t r i c a l conductivity of t r a n s i t i o n metal fluorosulfate 29 / \ complexes. Milne examined the conductivity of t e t r a k i s ( a c e t o n i t r i l e ; complexes of copper(l), c o p p e r ( l l ) , and z i n c ( l l ) fluorosulfates i n aceto-n i t r i l e (AN) solutions. He found that the copper(ll) and z i n c ( l l ) complexes have molar conductances much lower than those of 1: 1 e l e c t r o l y t e s . In contrast, the molar conductances of the corresponding complexes of .'manganese ( i i ) , c o b a l t ( l l ) , n i c k e l ( l l ) , c o p p e r ( l l ) , z i n c ( l l ) , and cadmium perchlo-r a t e s 1 0 2 ' 1 0 3 and copper(ll) t e t r a f l u o r o b o r a t e 1 0 ^ ' 1 ° ^ show normal 2 : 1 e l e c t r o l y t e behaviour. 29 Milne ascribes the low conductances of the copper(ll) and zinc ( i i ) fluorosulfate complexes as a r i s i n g from strong association of ions i n solution at the concentrations used |(0.1+5 - 1 . 5 ) x 1 0 _ 3M f o r Cu ( l l ) and (1 - 20) x 1 0 - 3M f o r Zn(ll)]. Very few e l e c t r i c a l conductivity studies have been made on 35 .51 8©. pyridine complexes of t r a n s i t i o n metal s a l t s . ?5 ;5'3"'-5 These measurements are very l i m i t e d i n scope, since none were taken over a concentration range. Furthermore, there i s very l i t t l e consistency i n either the concentration of the solution used, or the temperature at which the measurement was made. The molar conductances are also reported i n up to f i v e d i f f e r e n t solvents. In t h i s work, conductance measurements were made i n a c e t o n i t r i l e and dichloromethane solutions and i n pyridine where possible. Conditions _3 were standardized so that a l l measurements were made with 1 . 0 0 x 10 molar solutions at 25.00°C. 93 When choosing the optimum solvent f o r conductivity work, the c r i t e r i a which are most relevant are the d i e l e c t r i c constant, v i s c o s i t y , s p e c i f i c conductance, ease of p u r i f i c a t i o n , and the donor capacity of the solvent towards metal ions. Some properties of non-aqueous solvents relevant to e l e c t r i c a l conductivity work are given i n Table XXI. Solvent D i e l e c t r i c Constant V i s c o s i t y (cP) Donor Number A c e t o n i t r i l e Pyridine Dichloromethane Nitromethane Nitrobenzene Dimethylformamide Acetone 3 5 . 9 5 1 2 . 3 9 . 1 a 3 5 . 9 3 4 . 8 2 3 6 . 1 2 0 . 7 0 . 3 4 1 2 0 . 9 7 4 ' 0 . 3 9 3 * 0 . 6 2 0 2 . 0 3 a 0 . 7 9 6 O.316 a 11+.1 33-1 2 . 7 k.k 2 6 . 6 1 7 . 0 a = measured at 20 C. measured at 30 C. Table XXI - Some Properties of Non-Aqueous Solvents. The values f o r the d i e l e c t r i c constant and v i s c o s i t y are measured at 25°C unless otherwise specified. The Donor Number i s defined as the numerical quantity of the - AH^ g-^ Q^  value f o r the donor-accept or reaction: D ( d i s s ) + SbClw,. \ ^  D.SbCl,-5(diss) 5 9h 106 where D i s the donor solvent, and i s considered by Gutmann to be a semi-quantitative measure of solute-solvent interactions. A solvent with a high d i e l e c t r i c constant and a low v i s c o s i t y i s preferred f o r conductivity purposes, and on t h i s basis, a c e t o n i t r i l e was chosen f o r t h i s work. Use of t h i s solvent also permits comparison with the 29 work of Milne on the a c e t o n i t r i l e complexes of copper(l), eopper(ll), and z i n c ( l l ) fluorosulfates. While pyridine and dichloromethane are poor solvents f o r conduc-t i v i t y purposes, measurements of the molar conductances i n these solvents were included when t h i s i s permitted by the s o l u b i l i t y of the complex. This allowed a check on the anion d i s s o c i a t i o n e q u i l i b r i a i n these solvents since infrared, electron paramagnetic resonance and electronic spectra were also obtained with these solvents. The v a r i a t i o n of the conductance of copper(ll) tetrakis(pyridine) bi s f l u o r o s u l f a t e i n a c e t o n i t r i l e solution over the concentration range (0.1 - 12) x 10 3 moles 1 ^, at 2£°C was determined, and the data are given i n Table XXII. I f Cu(py)^(S0 3F)2 ionizes as a 2 : 1 e l e c t r o l y t e , then i t s equiva-lent weight w i l l be h a l f the formula weight. Making t h i s assumption, the equivalent conductance ( A g) and the equivalent concentration, C g were calculated, and the phoreogram (curve of A g plotted against square root of equivalent concentration) was plotted. This i s shown i n Figure 1£. In order to obtain the l i m i t i n g equivalent conductance, the 107 Shedlovsky's method of extrapolation was used. For t h i s purpose, values of A defined as-o • _ A-:+ A Jl (3.1) 0 _ 1 - B -n 96 Concentration (moles 1 x 10 ) Molar Conductance A m(ohm mole cm ) 11 .871 1 8 5 . 9 9-91+7 1 9 3 . 4 4-974 2 2 3 . 2 3-971+ 2 3 4 . 4 2 . 5 1 9 2 5 0 . 5 1 . 2 5 9 2 7 6 . 5 O .969 2 8 5 . 9 0 . 8 1 8 2 9 7 . 1 0 . 8 0 6 3 0 2 . 0 O .630 3 1 3 . 5 0 . 5 9 9 3 1 2 . 6 0.1+03 3 1 9 . 0 0 . 3 1 5 3 2 6 . 7 0.21+2 3 3 1 . 2 0 . 1 1 4 3 4 9 . 6 Table XXII - Conductance of Cu(-py). (£>0_F)_ i n A c e t o n i t r i l e Solution at 25°C. were evaluated from the experimental data and plotted against the io n i c strength, I , of the solution. This equation arises simply from a rearrangement of the Debye-Huckel-Onsager l i m i t i n g law which i s more usually found i n the form: 97 A = A Q - (A + B A o ) V T ( 3 - 2 ) Fe( |z | + |z_| ) / 87rNe2 X * where A 6 TrrjC x 1 0 " VlOOOekT e2co( |z +z_| ) / 87rHe2 ^ * 6ekT VlOOOekT 2 q and q 1 +V<i I V - I A o (|z+| + |z_|) (|z_| X Q + + | z + | X O " ) The symbols have the usual meanings: F = Faraday constant; e = electron charge; C = v e l o c i t y of l i g h t ; IT = Avogadro's number; e = d i e l e c t r i c constant; k = Boltzmann constant; 77 = v i s c o s i t y c o e f f i c i e n t ; T = temperature i n °K; z + , z_ = charges of cation and anion, respectively; Ar, X Q + , X. = l i m i t i n g conductances of the s a l t , the cation and the anion, respectively. 108 When the values of the physical constants are substituted into the expressions f o r A and B, the expressions reduce to the general forms: A = U 1 . 2 U 5 3 (|zj + IzJ) q. 1 +Vq 2 . 8 0 1 3 5 7 x 10 |z a j 98 For the bivalent copper(ll) fluorosulfate complex i n a c e t o n i t r i l e solution at 25°C, the constant A i s 3 5 0 . 2 7 8 , while B = 5.01+866 1 + Vq. where q - _o_ 3(2X 0" + X G +) The values of 3 5 . 9 5 and. 0.0031+12 poise were assumed f o r the d i e l e c t r i c 109 constant and the v i s c o s i t y c o e f f i c i e n t of a c e t o n i t r i l e , respectively. The l i m i t i n g equivalent conductance, X q , of the fluorosulfate ion i n 29 a c e t o n i t r i l e was estimated by extrapolating Milne's phoreogram on Et^NSO^F to i n f i n i t e d i l u t i o n . This procedure gave a l i m i t i n g conductance -1 -1 2 of 182 ohm equiv cm f o r the s a l t . Since the l i m i t i n g conductance of + 110 - 1 - 1 2 the Et^N ion i n a c e t o n i t r i l e i s known to be 8 5 . 0 5 ohm equiv cm , then, applying Kohlraush's law of the independent migration of ions, the l i m i t i n g conductance of the fluorosulfate ion i n a c e t o n i t r i l e was determined -1 -1 2 to be 97 ohm equiv cm . In c a l c u l a t i n g A q i n the f i r s t approximation of the Shedlovsky function, the value of 188 was assumed for A , as found by simple extra-polation of the phoreogram of Cu(py)^(SO^F^. In the second approximation, the value of A , found i n the f i r s t approximation was used i n the calcula-t i o n of A Q . Further approximations were unnecessary since the functions converge rapidly. The l i m i t i n g equivalent conductance thus determined f o r —1 — 1 2 Cu(py)^(S0 3F)2 i s 185 ohm equiv cm . The Shedlovsky plot used i n the evaluation of i s shown i n Figure 1 6 . 99 100 The Shedlovsky plot shows positive deviatons from the Onsager l i m i t i n g law at concentrations higher than approximately 1 x 10 M^. Such behaviour i s s i m i l a r to that observed for the divalent f i r s t row t r a n s i t i o n metal perchlorates i n a c e t o n i t r i l e solution where the s a l t s are f u l l y dissociated to give the |ML ^J 2 + ions (M = Mn, Co, N i , Cu, Zn, Cd). 107 As Shedlovsky has shown that application of the Onsager l i m i t i n g law does not give constant A q values even i n v e r y l d i l u t e solutions, his function i s tantamount to an empirical modification of the l i m i t i n g law to the form: . • = A + A V l _ T A o —TT + 01 ( 3 . 3 ) 1 - B V l where the slope of the A Q vs. I l i n e i s represented by C. However, the l i n e a r relationship between A q and I of equation ( 3 . 3 ) does not hold f o r Cu(py)^(S0 3F)2 (Figure 16), and the l i n e passes through a minimum at approxi--3 mately 1 .5 x 10 M with a positive deviation below t h i s concentration. 107 Shedlovsky has also observed non-linear behaviour at low concentration f o r potassium n i t r a t e , chlorates, iodates, and some thallium s a l t s i n aqueous solution. This deviation he attributes to ion association i n solution. Since the Shedlovsky function i s derived empirically, the reason why there should be a minimum i n the A q vs. I plot i s not understood. The / 102 Shedlovsky plot of Libus and S t r z e l e c k i f o r the t r a n s i t i o n metal perch-lorates show a positive deviation above 2 x 10 concentration, and a horizontal section down to 1 x 10 which they extrapolated to obtain the l i m i t i n g conductance. 101 102 So f a r , a l l data f o r CuXpy^SO^F^ have been analyzed by assum-ing that the complex ionises as a 2 : 1 e l e c t r o l y t e i n a c e t o n i t r i l e solution, but the v a l i d i t y of t h i s assumption may be e a s i l y v e r i f i e d by a determination of the slope of the l i n e a r portion of the phoreogram. According to the Onsager l i m i t i n g law A = - (A + B A Q ) y/l, the equivalent conductance varies l i n e a r l y with the square root of the i o n i c strength i n d i l u t e solutions of strong e l e c t r o l y t e s , and the slope of t h i s l i n e i s determined 111 l a r g e l y by the ion type and the solvent. A plot of ( A - A.) v s « \/1 (Figure 17) therefore y i e l d s a l i n e whose slope i s (A + B A Q ) . The slopes obtained from the experimental data, and calculated from the Onsager law by considering Cu(py)^(S0 3F)2 as a 1 : 1 or a 2 : 1 electrolyte are l i s t e d i n Table XXIII. Calculated Experimental 1 : 1 e l e c t r o l y t e 507 251+5 2 : 1 e l e c t r o l y t e 596 726 Table XXIII - Onsager Slopes f o r Cufpy^fSO^F)^ i n A c e t o n i t r i l e at 25°C The agreement between the experimental slope and the slope calculated from the Onsager law i s good evidence that Cu(py)^(S0 3F)2 does indeed behave as a strong 2 : 1 electrolyte i n a c e t o n i t r i l e solution. The difference between the experimental and calculated slopes arises i n part 103 from the approximations inherent i n the Onsager l i m i t i n g law, and also from the presence of some ion association which would cause a negative deviation from the Onsager l i m i t i n g tangent. 112 Harkness and Daggett i n studying the e l e c t r i c a l conductivity of tetra-n-alkylammonium s a l t s i n a c e t o n i t r i l e also observed that the experimental slopes of the phoreograms were greater than theoretical and that the difference between them was greatest i n those s a l t s where the degree of association was also highest. The occurrence of some ion association i s not unexpected i n view of the low d i e l e c t r i c constant of a c e t o n i t r i l e . • 102 103 Libus and St r z e l e c k i ' observed that the divalent t r a n s i t i o n metal perchlorates dissociate anions to give the [ M ( A N ) ^ ] 2 + and C10^ ions i n a c e t o n i t r i l e solution. The close s i m i l a r i t y of the conductivity data f o r Cu(py)^(S0.jF)2 i n a c e t o n i t r i l e solution suggests that the fluorosulfate complex ionises i n a l i k e manner to give a hexacoordinated copper(ll) complex ion. This i s also consistent with infrared r e s u l t s i n t h i s solvent which shows the presence of io n i c fluorosulfate ions, and no evidence of disso-ciated pyridine. Thus, the main species i n a c e t o n i t r i l e solution are probably [Cu(py)^(AlT) 2] and SO^F" ions.. Since the l i m i t i n g equivalent -1 -1 2 conductance A Q, of the complex i s found to be 18£ ohm equiv cm , i t follows that the l i m i t i n g equivalent conductance, X q + , of the cation i s - 1 - 1 2 — -1 -1 2 88 ohm equiv cm , i f X q i s assumed to be 97 ohm equiv" cm for the SO^F anion. This X Q + value f o r the [cu(py)^(AlT) 2] 2 + cation i s i n good —1 — 1 2 r i2+ 102 agreement with the value of 94.8 ohm" equiv cm f o r [Cu(AN)gJ . The lower value f o r [Cu(py)^(AN)^ + i s expected because the larger radius of the pyridine groups r e l a t i v e to that of a c e t o n i t r i l e , should lower the mob-i l i t y of the complex ion somewhat. The r e s u l t s obtained i n t h i s work, therefore, indicate that Cu(py)^(S0 3F) 2 behaves as a strong 2 : 1 ele c t r o l y t e i n a c e t o n i t r i l e solution as we would expect. In contrast to t h i s , however, M i l n e 2 ^ finds 104 unusually low conductances f o r copper(ll) and z i n c ( l l ) t e t r a k i s ( a c e t o n i t r i l e ) b i s f l u o r o s u l f a t e s , and he suggests that the complexes are strongly associated i n a c e t o n i t r i l e solution. I t i s d i f f i c u l t to understand why ion association i s not s i m i l a r l y important f o r the corresponding perchlorate and tetrafluoro-borate complexes which show normal 2 : 1 ele c t r o l y t e behaviour i n a c e t o n i t r i l e solution. Furthermore, Milne r a t i o n a l i s e s that the low conductance of C U X A H ^ ^ S O ^ F ^ must be due to the dipolar nature of the fluorosulfate anion which would favour pair formation. I f t h i s i s , i n f a c t , the true explana-t i o n of the low Ajyj values, then we should also expect strong ion associa-t i o n to be a general feature of fluorosulfate compounds i n a c e t o n i t r i l e solutions. This i s not observed, since Cu(py)^(S0 3F)2 shows sim i l a r conduc-t i v i t i e s to the corresponding perchlorate and tetrafluoroborate complexes for which ion association i s a minor consideration. Milne's arguments based on the importance of ion pa i r i n g therefore appear to be quite untenable. In order to compare the conductivities of complexes sim i l a r to Cu(py)^(lS0 3F)2, the molar conductances of a series of copper(ll) t e t r a k i s (pyridine) complexes were measured at 1 x 10 _ 3M concentration at 2J?0C i n a c e t o n i t r i l e , dichloromethane and pyridine. The re s u l t s are presented i n Table XXIV". The conductance of Bu^MBr i s included to indicate the values expected f o r a t y p i c a l 1 : 1 e l e c t r o l y t e . The re s u l t s show that i n a c e t o n i t r i l e solution, the conductances of the copper(ll) complexes can be r e a d i l y divided into two classes. The f i r s t class, comprising the tetrafluoroborate, perchlorate, and flu o r o s u l -fate complexes, has values which are t y p i c a l of strong 2 : 1 electro-? • 113 l y t e s . The p - t o l y l s u l f a t e , n i t r a t e , and trifluoroacetate complexes f a l l i n the other class, and have A ^ values lower than those expected for 1 : 1 105 e l e c t r o l y t e s . The range of molar conductances may arise from several factors, such as varying degrees of anion di s s o c i a t i o n , d i f f e r i n g anion m o b i l i t i e s , or ion association i n solution. Complexes —1 —1 2 Molar Conductance, (ohm mole cm ) CH^CN CH"2C12 Cu(py)^(BF u) 2 321 0 . 6 5 Cu(py) u(C10 u) 2 317 1 . 0 3 Cu(py) I t(S0 3F) 2 . 285 0 . 8 3 Cu(py) i +(CH 3C 6H i +S0 3) 2 76 0 . 3 1 1.1+ Cu(py) u(N0 3) 2 58 0 . 0 7 Cu(py) i +(CF 3C00) 2 11 0 . 0 2 2 . 0 Zn(py)^(S0 3F) 2 259 1 . 0 9 ~ 6 8 a Ni(py ) i | ( s o 3F) 2 221+ 9 ~ 7 5 b Bu^NBr 150 2 0 . 3 3 3 - 5 a = ~ 1 x 1 0~Sl. b = ~ 2 x 10 _ i 4M Table XXIV - Molar Conductances of Tetrakis(pyridine) Complexes. 106 I t i s u n l i k e l y that the anion m o b i l i t i e s are mainly responsible for the v a r i a t i o n of values, i n view of the wide v a r i a t i o n observed. To confirm t h i s point, the l i m i t i n g equivalent conductances of several anions i n a c e t o n i t r i l e at 2£°C are l i s t e d i n Table XXV. A l l the anions included i n t h i s conductance study, except for the p - t o l y l s u l f a t e anion, - 1 - 1 2 have X Q values which agree within 11 ohm equiv cm . The l i m i t i n g conductance of the p - t o l y l s u l f a t e anion has not been determined i n acetoni-t r i l e solution, but i t should l i e between those of the methylsulfate and the 2,5-dichlorobenzylsulfate anions. Thus, the small difference i n the molar conductances of the tetrafluoroborate and the perchlorate complexes may be due simply to the difference i n m o b i l i t i e s of the anions. However, t h i s w i l l not account for the large spread of values i n the second class of complexes. Consideration of m o b i l i t i e s w i l l have i t s greatest effect on the p - t o l y l s u l f a t e complex, but i t s t i l l w i l l not explain the wide gap between i t s e l f and the fluorosulfate complex. Hathaway and c o w o r k e r s 1 0 ^ ' 1 1 9 ' 1 2 0 have observed that copper(ll) n i t r a t e behaves as a very weak ele c t r o l y t e i n a c e t o n i t r i l e solution, i n close s i m i l a r i t y with our r e s u l t s f o r Cu(py)^(N0 3)2. Their r e s u l t s from u l t r a - v i o l e t absorption spectra, molecular weight, and conductivity measure-ments i n a c e t o n i t r i l e solution are consistent with each other i f i t i s assumed that stepwise dissociation: Cu(N0 3) 2 Cu(N0^) + + N0 3" C u 2 + + 2N03" _3 can occur. At 1 x 10 M concentration, t h e i r observations can be correlated on the basis of a f i r s t stage di s s o c i a t i o n i n which a n i t r a t e ion i s coval-ently bonded rather than present i n an ion-pair. 107 Limiting Equivalent Conductance, Anion . - / , -1 . - 1 2>. Reference X q (ohm equiv cm ; BF^~ 1 0 8 . l+a C10^" 1 0 3 - 7 115 SO^P" ~ 9 7 b N0 3" 106.1+ 111+ CP 3COO" 9 8 . 5 b 116 CH 3S0 3" 8 9 ° 117 CI CI SO," 7 0 b 117 PF 6 101+.2+ 111+ SCN" 113-1+ 118 calculated from Cu(A\tl)^X assuming X o~(C10^ ) = 1 0 3 - 7 ohm equiv cm . b calculated from Et^lT + s a l t assuming X o +(Et^N +) = 80 .05 ohm ^  equiv ^ cm2. C calculated from Na + s a l t assuming X Q +(Na +) = 7 6 . 9 ohm ^  equiv ^ cm2. Table XXV - Limiting Anion Conductances i n A c e t o n i t r i l e at 25°C. 108 From these r e s u l t s , then, we may explain the low conductance of Cu(py)^(N0 3) 2 as also a r i s i n g from a f i r s t - s t a g e anion dis s o c i a t i o n rather than a predominance of ion association. Thus, considering the series of copper(ll) tetrakis(pyridine) complexes, the dominant effect causing the wide v a r i a t i o n i n conductances i s the d i s s o c i a t i o n e q u i l i b r i a of the type: AN AN Cu(py) uX 2, 1 [cu(py) u(AN)] + + X~ [cu(py) u(AN) 2] 2 + + 2X_ The molar conductances should therefore show a v a r i a t i o n p a r a l l e l to the degree of d i s s o c i a t i o n of the anion i n the complex. Since the degree of d i s s o c i a t i o n i s d i r e c t l y related to the strength of coordination of the anions, we can arrange the anions i n order of increasing ligand strength from the conductivity data i n a c e t o n i t r i l e solution: BP, " < CIO," < S0.F" < CH-C^ H, SO ~ < NO ~ < CF-C06" 4 4 3 3 M 3 3 3 and f o r thenfluorosulfate complexes: Cu(l l ) < Z n ( l l ) < N i ( l l ) The copper(ii) tetrakis(pyridine) tetrafluoroborate, perchlorate, and fluorosulfate complexes therefore give molar conductances which indicate that the anions are completely, or almost completely dissociated i n acetoni-t r i l e solution to give 2 : 1 e l e c t r o l y t e behaviour. In contrast, f i r s t -stage anion d i s s o c i a t i o n predominates for the p - t o l y l s u l f a t e , n i t r a t e , and 109 t r i f l u o r o a c e t a t e complexes, and they show behaviour intermediate between those of non-electrolytes and 1 : 1 e l e c t r o l y t e s ; f o r fluorosulfate complexes i n d i l u t e a c e t o n i t r i l e solutions, we may also 29 explain the low conductances observed by Milne f o r the a c e t o n i t r i l e complexes of copper(ll) and z i n c ( l l ) fluorosulfate by a f i r s t - s t a g e disso-c i a t i o n , e.g. Such a r e s u l t i s unexpected i n view of the almost complete dis s o c i a t i o n of fluorosulfate anions observed for the corresponding pyridine complexes. Ho explanation can be found why the fluorosulfate complexes do not show conductances s i m i l a r to the perchlorate or tetrafluoroborate complexes. The molar conductances of the tetrakis(pyridine) complexes were also measured i n dichloromethane solution. Due to the occurrence of extensive ion association even at low concentrations as a re s u l t of the low d i e l e c t r i c constant and poor solvating properties of the solvent, i t i s not surprising that Table XXIV should indicate very l i t t l e anion displacement with low complex conductances i n t h i s solvent. The species i n dichloromethane solution may thus be formulated as J c ^ p y ) ^ ^ . Even i n the presence of extensive ion pairing, i t i s worthy of note that the copper(ll) complexes also show e s s e n t i a l l y the same order as i n acetoni-t r i l e solution. This makes i t a l l the more l i k e l y that the ordering of the complexes i s due more to the strength of the anion coordination than the extent of ion pairing. Since we have shown that the effect of ion association i s minor 110 In pyridine solution, only the copper(ll) p - t o l y l s u l f a t e and trif l u o r o a c e t a t e complexes are s u f f i c i e n t l y soluble to allow accurate measurement of the molar conductances. The r e s u l t s indicate that at most, only r e l a t i v e l y small amounts of i o n i c species are present, so that there i s l i t t l e tendency to dissociate anions to give to J c u(py)gj 2 + species. On the other hand, Zn(py)^(S0 3F)2 and Ni(py)^(S0 3F)2 show extensive anion diss o c i a t i o n . While i t i s not possible to determine whether the SO^F anions have completely dissociated because of extensive ion p a i r i n g and the very l i m i t e d s o l u b i l i t y of the complexes, i t appears that these complexes show molar conductances intermediate between 1 : 1 and 2 : 1 electrolytes. 111 U.U ELECTRONIC SPECTROSCOPY J4.I4..I Theory of Copper(ll) Spectra 9 The copper(ll) ion possesses a 3d outer electron configuration and i s a simple system t h e o r e t i c a l l y , as i t can be considered to be a single 10 positive hole i n an otherwise f i l l e d d configuration. In an octahedral cubic ligand f i e l d , the f i v e - f o l d degenerate 3d-orbitals of the copper(ll) ion s p l i t into two kinds, the lower t_ o r b i t a l s (d.^* d x z , d. ) being three-fold degenerate, and the upper e^ o r b i t a l s (d x2 y2» 1 ^'w0~^°^ degenerate. The single unpaired electron could be i n either of the components of the e state. The Jahn-Teller g effect requires any non-linear system with a degenerate ground state to undergo such a d i s t o r t i o n as w i l l remove the degeneracy. For the copper(li) ion, the degeneracy may be removed by elongated tetragonal or rhombic dis t o r t i o n s of the octahedron which r e s u l t i n having the odd electron i n the dx2_y2 o r b i t a l . Compressed structures r e s u l t i n having the odd electron i n the d ^ o r b i t a l . Detailed calculations have shown that the elongated structures are usually more energetically favourable than the compressed 121 structures, consistent with t h e i r more frequent occurrence. A good c r i t e r i o n of the stereochemistry of the copper(ll) ion i n a complex i s therefore i t s electronic spectrum as t h i s i s determined by the r e l a t i v e ordering and the energies of the one electron energy l e v e l s . For the copper(li) ion, these are simply related to the sequence of the symmetry 2 representations derived from the single D Russell-Saunders term. This relationship i s i l l u s t r a t e d for the copper(ll) ion i n an elongated tetragonal octahedral stereochemistry i n Figure 18. However, the electronic spectra not 112 .dp p x^-y^ 2_, 3 v 2 2 t 2 „ E = e .b„ .a. .b. g g 2 g 1 g 1£ g 2g-- d 2 2 z^ 2 2 •d , d xz yz 1g 2 >2g 'g 2g. "E •2_ L . 1 2 . 2 2g, g 2g 1g 1g B„_ = J1g 4,2 1 2 g 2g 1g 1g J+ v 2 2 J e^.b .a. .b g 2g 1g 1g Figure 18 - The Relationship between the Spectroscopic Terms and the One-Electron Energy Configuration f o r a Copper(ii) Ion i n an  Elongated Tetragonal Octahedral Ligand F i e l d of D.. Symmetry. only depend upon stereochemistry, but f o r a given stereochemistry, they vary with the value of 10Dq, and with the extent of the tetragonal d i s t o r t i o n . The energy-level diagram f o r l i g a n d - f i e l d s of symmetry (Figure 18) would predict three t r a n s i t i o n s , but the precise order of the energy le v e l s i n t h i s symmetry i s uncertain. Depending upon the degree of tetragonal d i s t o r t i o n , three possible energy l e v e l sequences may a r i s e : (a) 2B. < V < V < 2E ' 1g N 1g v 2g g (b) 2B, < 2 B 0 ( V < 2E w 1g 2g 1g g 113 (c) \ < 2B„ < 2E < 2A, 1 g 2g g 1 g Case (a) w i l l correspond to a low tetragonal d i s t o r t i o n , and (c) to a large tetragonal d i s t o r t i o n which might be found i n true square-coplanar complexes. The infrared (section i | . . 2 . 2 ) and electron paramagnetic resonance (section I4..6) spectra, f o r the copper(ll) complexes studied, show no evidence for a square coplanar J c u ( p y ) J 2 + cation, and a l l the complexes involve an elongated tetragonal octahedral stereochemistry. Single-crystal electronic and electron paramagnetic resonance spectra of the tetragonal octahedral bis(ethylenediamine) and tetra-ammine copper(ll) complexes 8^'"' 2 2''' 2 3 have established the sequence of energy l e v e l s i n these complexes as: 2 2 2 B. < B 0 < E 1 g ^ 2 g s g 2 2 2 with A„ probably l y i n g between the B 0 and B„ l e v e l s . 1 g 2 g 1 g I4..I4..2 Results and Discussion f o r Copper(ll) Complexes The electronic spectra of the series of tetrakis(pyridine) copper(ll) complexes were obtained i n the s o l i d state from diffuse reflectance and nujol or hexachlorobutadiene mulls at room temperature. Spectra were also obtained from solutions i n dichloromethane, pyridine, and mixtures of these. The r e s u l t s are given i n Table XXVI. S o l i d State Electronic Spectra The diffuse reflectance and mull spectra of the pyridine complexes studied i n t h i s work consist of a single asymmetrical band envelope which Solid State Solution Diffuse Reflectance Mull CH 2C1 2 CH2C12<;.+ 10% py CH 2C1 2 + 20% py Pyridine 76,121+ 1 8 . 3 0 , I6 .70sh 1 9 . 1 , 1622-sh - - - -NO^ " 1 8 . 0 8 1 8 . 3 5 1 4 . 1 6 (81) M 6.61 =.(55) 1 6 . 7 2 ( 5 6 ) i n s o l . 1 7 . 7 6 - 1 7 . 7 9 ( 6 0 ) 1 7 . 7 9 ( 6 1 ) - ti 1 7 . 6 7 1 7 . 9 5 1 8 . 0 5 ( 5 3 ) 1 7 . 9 9 -FSO^" 1 7 . 0 4 1 7 . 2 4 1 7 . 2 7 ( 5 2 ) 1 7 . 1 2 ( 5 6 ) 1 7 . 1 2 ( 5 6 ) 1 5 . 9 2 C H 3 C 6 V ° 3 " 1 6 . 5 0 1 6 . 8 6 1 3 . 9 6 ( 7 6 ) 1 6 . 5 6 ( 5 2 ) - 16 . 5 3 ( 5 3 ) _ 3 0 , 7 7 ReO^ - 1 6 . 4 - - - -CF^COO" 1 5 . 5 3 1 5 . 7 7 1 4 . 9 9 ( 7 4 ) 1 4 . 6 8 ( 9 1 ) - 1 5 . 3 8 ( 6 4 ) molar ab s o r p t i v i t i e s , e , i n brackets. Table XXVT - V i s i b l e Spectral Data (kK) f o r Cu(py), X ? Complexes. 175 4 0 0 M U L L S P E C T R A 1 Cu(py) 4 ( N 0 3 ) 2 2 Cu(py) 4 ( B F 4 ) 2 3 Cu(py) 4 ( S 0 3 F ) 2 4 Cu(py) 4 ( C f ^ C 0 2 ) 2 6 0 0 8 0 0 1000 X ( n m ) Figure 19 116 spans the entire v i s i b l e portion of the spectrum. The mull spectra of several of the complexes at room temperature are shown i n Figure 1 9 . At temperatures close to that of l i q u i d nitrogen, i t has been shown that better resolution of the spectrum can be obtained. Thus, the room tempera-ture spectrum of Cu(py)^(PFg)2 shows a weak shoulder on the low energy side of the main absorption band, which on cooling to -196°C, i s resolved into a 76 separate band of medium i n t e n s i t y at 1 6 . 6 kK. A single asymmetrical absorption band i s frequently observed for copper(ll) compounds because the r e l a t i v e energies of the d-d t r a n s i t i o n s involved generally occur within 5 - 0 kK of each other. For elongated t e t r a -gonal octahedral complexes, an increase i n the degree of tetragonality raises 2 the energy of the state, and a s i t u a t i o n may arise i n which t h i s state 2 2 i s s u f f i c i e n t l y close to the E^ and 3^ states f o r the three t r a n s i t i o n s not to be resolved i n the spectrum, since the half-widths of these t r a n s i -125 tions can be as large as 3 * 0 kK« Hathaway8^ has observed that the ^B^g —** t r a n s i t i o n i s the most intense t r a n s i t i o n , and l i e s at the higher energies i n s i m i l a r tetragonal compounds. The band maxima of these compounds therefore probably represent 2 2 a close approximation to the B„ *• E t r a n s i t i o n with the remaining 1 g g t r a n s i t i o n s hidden on the low energy side of the band to give an asymmetrical band shape. The diffuse reflectance spectra of most of the complexes studied here have been reported previously, and good agreement i s found with the e a r l i e r r e s u l t s . Thus single absorption bands have been observed i n the s o l i d state spectra of tetrakis(pyridine) copper(ll) perchlorate,34>35J121J. tetrafluoroborate, 3^ nitrate,"''' and f l u o r o s u l f a t e . 3 0 Agambar and O r r e l l 8 2 117 report two bands i n the diffuse reflectance spectrum of the trifluoroacetate complex at 1 5 - 9 5 and I 4 . 8 O kK whereas we observe a single, very broad band with a maximum absorbance at 1 5 - 5 3 kK. 123 Hathaway and coworkers have suggested four c r i t e r i a f o r measur-ing an increase i n tetragonal d i s t o r t i o n , from tetragonal octahedral towards a square coplanar configuration. The greater the d i s t o r t i o n : 2 2 ( i ) the higher i s the energy of the B.. *• A., t r a n s i t i o n ; ' g ' g 2 2 ( i i ) the higher i s the energy of the IL *• E t r a n s i t i o n , ' c~> c~> [but to a lesser extent than i n (i)] ; ( i i i ) the shorter i s the in-plane copper-ligand bond length, R ; s ( i v ) the smaller i s the r a t i o T = Rg/R^ » where the out-of-plane Q copper-ligand bond length, RT involves a correction f o r the h tetragonal ligand atom present. A t o t a l absence of crystallographic data f o r the tetrakis(pyridine) copper(ll) complexes from s i n g l e - c r y s t a l x-ray d i f f r a c t i o n precludes the use of c r i t e r i a ( i i i ) and ( i v ) f o r comparing the tetragonality of these complexes. 2 2 However, we have a good estimate of the B„ > E t r a n s i t i o n 1g g from the band maxima of the complexes, and c r i t e r i o n ( i i ) permits a direct comparison of the degrees of tetragonal d i s t o r t i o n i n the complexes l i s t e d i n Table XXVI. By arranging the copper(ll) complexes i n order of decreasing band maxima, we obtain f o r the degree of tetragonal d i s t o r t i o n : PP/- > H0_ > CIO, > BP, ) PSO- > CH-.C.H, SO ~ > ReO, " > CF-,C00" 6 3 k k 3 3 6 4 3 h 3 Since the most strongly coordinating a x i a l ligands w i l l tend to coordinate to the copper atom and form tetragonally distorted octahedra i n 118 which the difference i n in-plane and out-of-plane f i e l d strengths w i l l be small, the reverse order w i l l also duplicate the coordination strength of the ligands: CF^COO" > ReO^" > CH^H^SC^" > FSC^" > BF^~ ) C10^~ > NC>3" > PFg" 126 Lever and Mantovani have made a systematic study of copper(ll) complexes of the general formula Cu(N-N) 2X 2 where N-N = ethylenediamine (en), 1,3-propylenediamine and N-substituted dimethyl- and diethylethylenediamines. They observe a straightltfine c o r r e l a t i o n between the square of the highest metal-nitrogen stretching v i b r a t i o n and the energy of the main electronic d-d band. They conclude that the position of a complex on the l i n e i s a measure of i t s tetragonal d i s t o r t i o n and of the in-plane bond strength. From t h e i r data on the ethylenediamine complexes, they obtain a spectro-chemical series: en > I " > Br" > CI ) CIO^" > SCN" > BF^~ > A g l 2 ~ With dif f e r e n t in-plane ligands, the order of the anions may vary, e.g. from asymmetric diethylethylenediamine (Et 2en) data, the series i s : Br" > C l " > BF^" > CIO^" > N0 3~ I f the energy of the main electronic d-d band i s used as the sole c r i t e r i o n of the tetragonal d i s t o r t i o n , e s s e n t i a l l y the same series i s obtained as from the correla t i o n with the M-N stretching v i b r a t i o n . Consid-ering the electronic spectrum only, a spectrochemical series f o r the Wavelength (nm) Figure 20 - Solution Electronic Spectra of Cu(py) X Complexes Wavelength ( nm) Figure 20 continued 121 ethylenediamine complexes i s obtained with the positions of Agl^ and BF^ reversed. The series remain •unchanged for the asym-Et^en complexes. I t i s clear then, that the position of the main electronic absorption band i s a -useful i n d i c a t i o n of the degree of tetragonal d i s t o r t i o n , though i t i s preferable to use i t s cor r e l a t i o n with the square of the highest metalsnitrogen stretching v i b r a t i o n when t h i s information i s available. Solution Electronic Spectra By comparison with the s o l i d state electronic spectra, Table XXVT shows the band maxima of the fluo r o s u l f a t e , perchlorate, and tetrafluoro-borate complexes to be l i t t l e changed i n dichloromethane solution. The band maxima of the n i t r a t e , p - t o l y l s u l f a t e , and trifluoroacetate on the other hand, show a r e l a t i v e l y large red s h i f t . The absorption spectra of the complexes i n solution are shown i n Figure 2 0 . The behaviour of each tetrakis(pyridine) copper(ll) complex i n solution w i l l now be considered i n turn. B i s ( f l u o r o s u l f a t o ) t e t r a k i s ( p y r i d i n e ) c o p p e r ( l l ) dissolves i n dichloromethane to give a purplish-blue solution. The close agreement of the band maxima of t h i s solution with those of the diffuse reflectance and mull spectra, indicates that the complex i s e s s e n t i a l l y unchanged i n solu-t i o n . There i s no further change i n band maximum i f the solution i s allowed to stand for f i v e days. The infrared spectrum also shows no evidence of dissociated pyridine even i n concentrated solutions (section I 4 . 2 . I ) . On addition of up to 25% pyridine to the solution, there i s only a small s h i f t to lower energy i n the electronic absorption band which may arise from some 122 substitution of pyridine ligands f o r fluorosulfate within the coordination sphere. As the frequency s h i f t i s small, such an effect i s minor and the species i n dichloromethane solution i s predominantly jcu(py)^(S0 3F) ?J. The complex i s s l i g h t l y soluble i n pyridine, r e s u l t i n g i n a l i g h t blue solution. The band maximum at 15•92 kK suggests that the fluorosulfate complex has dissociated anions i n t h i s solvent to give the j c u ( p y ) g j 2 + species. This interpretation i s consistent with the infrared spectrum of the solution which indicates i o n i c fluorosulfate groups with symmetry. In a c e t o n i t r i l e solution, a blue solution i s obtained with a band maximum at 16.81+ kK and a molar absorptivity e, of 60 1. mole, cm . On addition of 10% pyridine to the solution, the Diaximum s h i f t s to 1J.2lx kK with an increase i n e to 68 1. mole, cm , probably due to a suppression of any pyridine d i s s o c i a t i o n from the complex. The infrared spectra and e l e c t r i c a l conductivity both indicate complete fluorosulfate ion dissociation. Cu(py)^(C10^) 2 shows e s s e n t i a l l y no difference between the electronic spectrum of the dichloromethane solution and the s o l i d state spectrum. The band maxima and molar absorptivity are i n excellent agreement 35 with previously reported values. Addition of pyridine effects no change i n the band maximum. Cu(py)^(BF^) 2 also shows a simi l a r agreement between the solution and s o l i d state spectra. The perchlorate and tetrafluoroborate complexes therefore show no evidence of anion di s s o c i a t i o n , and the species i n solution are £cu(py)^(C10^) 2 J and |^Cu(py)^(BF^) 2J , respectively. The electronic spectra of Cu(py)^(CFC00) 2 indicate that the complex undergoes s i g n i f i c a n t changes when i t i s dissolved i n dichloromethane. The band maximum at 1i)..99 kK i n the solution spectrum may be considered to be due l a r g e l y to the bispyridine complex i n agreement with infrared 123 evidence. Admixture of pyridine suppresses the ligand d i s s o c i a t i o n some-what, hut not enough to bring a l l of the complex back to the t e t r a k i s (pyridine) form. In pure- pyridine solution, there appears to be very l i t t l e d i s s o c i a t i o n of the strongly coordinated trifluoroacetate group, and the complex remains as | c u(py)^(CF 3C00) 2 J . The band maximum for Cutpy^CH^CgH^SO^ shows a s h i f t of 2 . 9 0 kK to lower energy when the complex i s dissolved i n dichloromethane solvent. The e l e c t r i c a l conductivity of t h i s solution indicates l i t t l e d i s s o c i a t i o n of p - t o l y l s u l f a t e groups, and the infrared spectrum shows considerable d i s s o c i a t i o n of pyridine. Addition of 10% pyridine to the solution suppresses t h i s d i s s o c i a t i o n and brings the band maximum to w i t h i n 0 . 3 0 kK of that i n the mull spectrum. In pyridine solution, the agreement i s within 0 . 3 3 as evidence that the greater part of the complex i s s t i l l present as [ c u(py) 4(CH 3C 6H uS0 3) 2] . The electronic spectrum of Cu(py)^(N0 3) 2 i n dichloromethane solu-t i o n shows a s h i f t of 3 ' 9 2 kK to lower energy when compared with i t s mull 51 spectrum. Haendler and coworkers have reported a band maximum of "11+. 29 kK i n chloroform, s i m i l a r to the maximum at 11+.16 kK observed here i n dichloro-methane. Such a dramatic departure from the s o l i d state value i s most l i k e l y due to d i s s o c i a t i o n of pyridine to give the bispyridine complex with bidentate n i t r a t e groups. This interpretation i s consistent with the i n f r a -red spectra and e l e c t r i c a l conductivity which show extensive pyridine d i s s o c i a t i o n and no n i t r a t e ion dis s o c i a t i o n . Addition of up to 20% pyridine brings the band maximum i n closer agreement with the mull spectrum, though there i s s t i l l a difference of some 1 .6 kK between the solution and s o l i d state spectra. Thus, even i n the presence of such a large excess of free pyridine, the bispyridine complex 121+ does not completely revert back to the tetrakis(pyridine) complex. The species i n solution may be a mixture of the b i s - , t r i s - , and t e t r a k i s (pyridine) complexes. 1+.1+.3 Theory of N i c k e l ( l l ) Spectra g The n i c k e l ( l l ) ion has a 3& valence electron configuration which gives r i s e to the Russell-Saunders terms ( i n order of increasing energy) 3 1 3 1 1 F, D, P, G, S. In an octahedral c r y s t a l f i e l d , the degeneracy of the F term i s removed; the A Q term corresponds to the ground state (t„ ) (e ) <^g ^g g configuration i n either weak or strong f i e l d formalism. A Laporte forbidden d-d electronic t r a n s i t i o n , made possible through vibronic coupling, w i l l produce the excited states. The three spin-allowed electronic t r a n s i t i o n s available to the n i c k e l ( l i ) ion i n an octahedral c r y s t a l f i e l d are 125 Tetragonal d i s t o r t i o n a r i s i n g from a decrease of the ligand f i e l d strength along the z-axis causes a lowering of the molecular symmetry from 0^ to D J ^ J and the three spin-allowed excited electronic states are s p l i t , thus permitting s i x spin-allowed t r a n s i t i o n s . The correla t i o n diagram for n i c k e l ( l l ) i n 0^ and c r y s t a l f i e l d s i s shown i n Figure 21. Expressions r e l a t i n g t r a n s i t i o n energies to the r a d i a l parameters g Ds, Dt, Dq^ -y.' and Dq^ have been developed from c r y s t a l f i e l d theory f o r d 127 trans-ML^X^ type complexes. Ds and Dt describe the s p l i t t i n g i n the 3 three octahedral T states, while Dq^ and represent the out-of-plane and in-plane ligand f i e l d strengths, respectively. Equating the energy of the 3 ground state to zero, the energies of the f i r s t two excited states, assuming the absence of configurational i n t e r a c t i o n between le v e l s of the same symmetry, are to a f i r s t - o r d e r approximation, 1 2^ E ( 3 B 2 g ) = 1 0 D V E ( 3E^) = lODq^ - ^ D t E ( 3 A A G ) = I O D C L ^ + 12B - i+Ds - 5Dt E ( 3Eg) = 1 0 D q x y + 1 2 B + 2 D s " T^Dt 3 3 The B^ —>- 2g ^ ^ s i ^ i 0 1 1 ^ n t h i s approximation i s dependent only on the in-plane ligand f i e l d strength E q ^ . The s p l i t t i n g of the f i r s t excited state (^T2g> Qh) i s 3 - ^ " D t' w h i l e "tha't o f t h e s e c o n ( i (^ Tig' 0h) i s 6Ds - 5 A Dt. The signs of Ds and Dt are taken to be positive f o r complexes 129 with a x i a l elongation. Once Dq^. and Dt are obtained, c a l c u l a t i o n of the a x i a l f i e l d parameter, Dq^ i s possible from the expression: 'I 126 Dt = 4 ( D V - D q z ) fo r D j ^ molecules of the type trans-ML^X^. J4.J4.J4. Results and Discussion for N i c k e l ( l l ) Complexes S o l i d State Electronic Spectra The room temperature diffuse reflectance and mull spectra of N i ( p y ) ^ ( S 0 3 P ) 2 show structure which are t y p i c a l of octahedral complexes of the trans-NiL^Xg type. The spectra of the complex compared with that of N i ( p y ) ^ ( 0 1 0 ^ ) 2 are shown i n Figures 22 and 2 3 . The nujol mull spectrum of 3li 80 1 30 1 31 N i ( p y ) ^ ( 0 1 0 ^ ) 2 has been reported previously, ' both at room and l i q u i d nitrogen temperature. The r e s u l t s obtained here are i n excellent agreement with the e a r l i e r work. The s i m i l a r i t y between the spectra of the fluorosulfate and the perchlorate complexes indicates that Ni(py)^(SO^F^ i s distorted i n a s i m i l a r manner to the tetragonal octahedral Ni(py)^( ClO^^. ^ n e bands observed i n the mull spectrum of Ni(py)^(S0 3F)2_were therteffcoiae assigned by 1 31 comparison with Ni(py)^(ClO^^ according to Peeters . and these are given i n Table XXVII. Because of the l i m i t e d wavelength range of the spectrophotometer, only two band maxima were obtained i n the diffuse reflectance spectra. These occurred at 16 .26 and 26.I46 kK f o r N i ( p y ) ^ ( S 0 3 F ) 2 , and 16.67 and 2 6 . 8 2 kK for N i ( p y ) U ( C 1 0 U ) 2 . The r e s u l t s of c r y s t a l f i e l d calculations on several n i c k e l ( l l ) complexes at room temperature, performed by the procedure described i n section I4.l4.3j are given i n Table XXVIII. 127 Transition Ni(py)^(S0 3F) 2 Ni(py)^(C10^), \ V ( 3F) 8.1+8 7 . 2 2 1 1 . 7 6 sh 12.08 \ 0 » ) j 13 .61 sh 13.1+4 34s ) 3E* ( 3F) 16.1+2 1 6 . 7 2 ^ ( 3 P ) ^ 2 7 . 7 sh* ~2l+.l+ sh V ( 3P) 2 7 . 1 7 2 7 . 1 0 taken from diffuse reflectance spectra. Table XXVII - Room Temperature Electronic Spectral Data (kK) f o r  NifpyySQ-.Fk and M f p y ^ f C l O ^ as Kel-F Mulls. Anion, X Dt Ds BF^" 1290 90 686 680 ClO^" 1208 186 613 67I+ ReO. ~ 1190 1+58 1+18 601+ 4 S0 3F" 1176 519 375 51+7 Br" 1110 510 343 U05 CI" 1100 651+ 255 312 Table XXVIII - Spectrochemical Parameters f o r Ni(py)^X 2 Complexes  at Room Temperature. Diffuse Reflectance Spectra N i ( p y ) 4 ( C l 0 4 ) 2 Ni ( p y ) 4 ( S 0 3 F ) 2 3 5 0 4 0 0 4 5 0 5 0 0 5 5 0 6 0 0 Wavelength (nm) 6 5 0 7 0 0 7 5 0 Figure 22 - Room Temperature Diffuse Reflectance Snectra of Ni(pv) (SO^F) 2 and N i ( p y ) 4 ( C 1 0 4 ) 2 I r o Wavelength(ran) Figure 23 - Room Temperature Kel-F Mull Spectra of Ni (py) (SO F) and Ni(py) (CIO ) 130 The spectrochemical parameters f o r Ni(py)^(BF^) 2 were calculated 80 from the room temperature mull spectrum of Rosenthal and Drago. Since they did not report the position of the lowest energy t r a n s i t i o n , t h i s was taken as 6 . 9 0 kK from the spectrum at -196°C,"' 30 assuming a s h i f t to higher energy on warming to room temperature comparable i n magnitude to that of the perchlorate complex. ' The position of the »- 2 g " k r a n s i i ; : * - o n was taken as coinciding with that of the spin-forbidden t r a n s i t i o n s . The parameters f o r Ni(py)^(Re0jp2 were obtained from the spectral data of M a y f i e l d . 3 0 The ^B-jg *" ~^2g " ^ ^ s i ^ i 0 1 1 w a s assumed to coincide with the band at 11+.00 kK. The c r y s t a l f i e l d parameters f o r the bromide and chloride complexes were obtained from the diffuse reflectance spectra of Goodgame 132 and coworkers. In c a l c u l a t i n g the out-of-plane f i e l d strengths, i t i s necessary to use Dt, as determined d i r e c t l y from the s p l i t t i n g of the f i r s t spin-allowed absorption band. Since the spin-allowed tra n s i t i o n s concerned s p l i t into two components as a direct consequence of the tetragonal component to the ligand f i e l d , the degree of s p l i t t i n g i s proportional to t h i s tetragonal component and Ito Dt. Therefore, used i n a r e l a t i v e sense, both Dt and Dq z derived therefrom, can be considered to r e f l e c t the strength of the out-of-plane ligand f i e l d . Comparison of the r e s u l t s of the calculations performed here with those taking into account configurational i n t e r a c t i o n (Cl) shows that the in c l u s i o n of CI between A 0 and between E l e v e l s does not greatly affect 2 g g the evaluation of Dt, and hence of Dq , but may materially a l t e r the other parameters. Hence the Dq z data given i n Table XXVTII where CI was ignored, may be accepted as approximately v a l i d . A further approximation of the c r y s t a l f i e l d parameters arise from the use of the room temperature spectral data. At the ambient temperature, the bands are broader than at l i q u i d 131 nitrogen temperature, and some resolution of the trans i t i o n s i s l o s t . In 3 3 a pa r t i c u l a r , the >• Ag hand i s usually not resolved, and overlaps with the spin-forbidden t r a n s i t i o n s . Making the assumption that the 3 3 a JTS>. J k n band coincides with that of the spin-forbidden tr a n s i t i o n s 1g 2g could lead to appreciable error i n the Ds parameter. However, i f the spectra of the n i c k e l ( l l ) complexes are analyzed under the same conditions, i . e . from room temperature data and excluding CI, then comparisons between the complexes are s t i l l v a l i d , since only the r e l a t i v e magnitudes of the parameters are considered, although the actual values may be appreciably i n error. The changes i n the c r y s t a l f i e l d parameters from room and low temperature spectral data, and with or without the inc l u s i o n of a l l config-urational i n t e r a c t i o n are shown f o r Ni(py)i (CIO. )_ i n Table XXIX. Without CI With CI a 6" room temp. low temp. low temp. Dq^. 12^ 8 1292 1289 Dq 186 282 2+10 z Dt 613 577 500 Ds 672+ 750 930 a. Spectral data from r e f . 13©. b. Parameters from r e f . 13.0. Table XXIX - Crystal F i e l d Parameters (cm ) for Ni(py)^(CIO^) g. 1 3 2 I f we arrange the Dq^ values for several Ni(py)^X£ complexes obtained from the electronic spectra at -196°C and calculated with i n c l u s i o n 30 129 of CI, we obtain the spectrochemical series: C l " ) ReO^" ) Br" > C10^" ) BF^~ _-| with the corresponding Dq values of 6 8 0 , 6 6 8 , 6 0 0 , 1 + 1 0 , and 3 6 0 cm , z respectively. A comparison of the Dt and Dq^ values i n Table XXVIII from the room temperature electronic spectra therefore indicates that the fluorosulfate ion coordinates more strongly than the perchlorate and tetrafluoroborate ions, and i s s i m i l a r i n coordinating a b i l i t y to the chloride, bromide, or perrhenate anions. Solution Electronic Spectra Ni(py)^(S0jF)2 dissolves i n dichloromethane to give a green solution which shows considerable change i n band positions r e l a t i v e to those i n the s o l i d state spectrum. The e l e c t r i c a l conductivity of the solution indicates s i g n i f i c a n t d i s s o c i a t i o n of fluorosulfate ions, which i s probably the main cause of the large s h i f t i n the band maxima. No evidence of pyridine disso-c i a t i o n i s observed i n the infrared spectrum of the solution. On addition of up to 3 0 % pyridine to the solution, there i s a steady s h i f t of the band maxima to higher energy concomitant with an increase i n .the molar absorpti-v i t i e s of the bands. As the proportion of pyridine i n the solution increases, the absorption band at ~ 1 6 . 0 kK becomes narrower, and the shoulders at 1 1 . 7 6 and 13.61 kK decrease i n i n t e n s i t y so that the spectrum has a closer resemblance to that of a regular octahedral n i c k e l ( l l ) complex. 133 The changes i n hand maxima and molar a b s o r p t i v i t i e s on addition of pyridine to the dichloromethane solution are given i n Table XXX. The corresponding bands i n the room temperature mull spectrum are included f o r comparison. "1 \ "3 Mull 8.1+8 16.1+2 2 7 . 1 7 CH 2C1 2 8 . 8 5 (1+) 1 5 . 0 2 ( 7 ) 2 5 . 2 5 (11+) CH 2C1 2 + 5% pyridine 9 . 1 3 (1+) 1 5 - 9 7 ( 7 ) 26.01+ (11+.5) CH 2C1 2 + 10% pyridine 9 . 4 5 ( 5 ) 1 6 . 1 8 ( 9 ) 26.18 (18) CH 2C1 2 + 20% pyridine 9 . 5 7 ( 5 ) 1 6 . 2 3 ( 9 ) 2 6 . 3 2 ( 1 8 ) CH 2C1 2 + 30% pyridine 9 - 7 6 16.1+5 -Pyridine - 16 .53 2 6 . 6 7 Table XXX - Band Maxima (kK) for Ni(py), (S0.P) „ i n mixed CH„C1„-Pyridine Solution. The s h i f t s of the band maxima and the change i n appearance of the spectrum are consistent with an increased d i s s o c i a t i o n of the fluorosulfate ion to give the | N i ( p y ) g J 2 + cation as the proportion of pyridine i n the solution i s increased. The changes i n the electronic spectrum may be seen i n Figure 21+. -p •H > •H •P ft fH O CQ 15H 10 H 5 H 0 300 C H 2 C I 2 CH Cl 2+25% Pyridine 500 700 9 0 0 1100 1300 1500 ( nm) Figure 24 - Electronic Spectra of Nifpy) (SC^F) i n CH 2C1 2 and CH 2C1 2 + 25% Pyridine Solution 135 N i c k e l ( l l ) t e t r a k i s ( p y r i d i n e ) b i s f l u o r o s u l f a t e i s sparingly soluble i n pure pyridine. The lowest energy band i n the electronic spectrum of the blue solution i s too weak and broad to obtain a good estimate of i t s energy. 80 The band maxima for Ni(py)^(BF^) 2 i n pyridine solution have been reported at 9»57> 16.1+77 and 26.7I+ kK, i n good agreement with those obtained here f o r 80 the Ni(py)| +($0n^})^ 2 complex. Rosenthal and Drago report a 10Dq for |lTi( p y ) g J 2 + of 9«85 kK obtained from the lowest energy band of the perchlo-rate and tetrafluoroborate complexes i n nitromethane solution containing a large excess of pyridine. The electronic spectrum of Ni(py)^(S0 3F) 2 i n a c e t o n i t r i l e solution has the appearance of a t y p i c a l tetragonal octahedral n i c k e l ( i i ) complex. Absorption band maxima were observed at 10.93> 17«67 and 2 8 . 1 7 kK, with shoulders at 25.0 kK and 1 3 . 3 kK. The s p l i t t i n g of the band at 17.67 kK i s not resolved, but i t does show a r e l a t i v e l y large band width with a pro-nounced asymmetry toward the low energy side of the band. The band maxima observed here are i n excellent agreement with those reported by Rosenthal 80 and Drago f o r Ni(py)^(C10^) 2 i n a c e t o n i t r i l e . They observe an average stoichiometry of | l T i ( p y ) 2 ( A N ) J 2 + f o r the perchlorate complex i n t h i s solution. The agreement of the electronic spectra f o r the two complexes suggests ;a s i m i l a r stoichiometry f o r Ni(py)^(SC> 3F) 2 i n a c e t o n i t r i l e solution. 136 k.S MAGNETIC SUSCEPTIBILITY I4..5.I Theory of Magnetic S u s c e p t i b i l i t y Under the effect of an octahedral c r y s t a l f i e l d , some of the 2 o r b i t a l degeneracy of the D ground term of the copper(li) ion i s l i f t e d , 2 and i t i s s p l i t to give a lower energy doublet E term, and an upper t r i p l e t 2 2 T_ term (Figure 1 8 ) . The E ground term corresponds to an approximate g 6 3 electron configuration of e^. An inspection of the nature of the angular part of the wave functions for the e^ o r b i t a l s (d x 2 y 2 ' ^ z 2 ) indicates that no simple r o t a t i o n process can transform one o r b i t a l into the other, as they d i f f e r i n shape. Although they are degenerate from an energy viewpoint, there i s no.' o r b i t a l momentum associated with the e group, since the presence of o r b i t a l momentum i n an & ion may be associated with the a b i l i t y to interchange one o r b i t a l into another 6 3 by rotation. Thus, for a e^ electron configuration with a "positive" hole i n an e o r b i t a l , no o r b i t a l contribution i s expected i n the magnetic moment. The magnetic moment of octahedral copper(ll) complexes should, therefore, be close to the spin-only value of 1 . 7 3 - However, the moments obtained experimentally i n the absenceL of any antiferromagnetic coupling are i n the range 1 . 8 - 2 . 0 , appreciably above the spin-only value. S i m i l a r l y , i n an octahedral complex, the n i c k e l ( l l ) ion has a 3 A~ ground term, for which no o r b i t a l contribution to the magnetic moment 2g i s expected. The actual range of moments found i n spin-free n i c k e l ( l l ) com-plexes i s 2 . 9 - 3 - 5 > as compared with the spin-only value of 2 . 8 3 . The increase above the spin-only values arises from perturbation by spin-orbit coupling, causing the "mixing-in" of some of the o r b i t a l l y 137 2 degenerate excited state. As a consequence, the separation of the E g 2 3 3 and Tg terms f o r co p p e r ( l l ) , and the ^ and T^g terms f o r n i c k e l ( l l ) on the basis of t h e i r d i f f e r i n g o r b i t a l angular momentum, i s not e n t i r e l y v a l i d . The ground terms may therefore r e t a i n some o r b i t a l angular momentum, which i s 2 3 equivalent to saying that they are not pure E g or terms. There i s also a contribution to the magnetic s u s c e p t i b i l i t y of E or A ground terms from the second-order Zeeman effect which mixes into -the ground state some of the character of a higher state, causing a lowering of the ground state. When the separation of the inte r a c t i n g l e v e l s i s much greater than kT, the inte r a c t i o n i s independent of temperature and i s there-fore referred to as temperature independent paramagnetism ( t i p ) . For octahe-2 2 dra l copper(ll) complexes with E ground terms, t i p = 1+1T/3 /lODq, and for 3 2 octahedral n i c k e l ( l l ) complexes with S r o u n ^ terms, t i p = 8N/3 /lODq, where N = Avogadro's number, and /3 = Bohr magneton. In order to f u l l y account f o r the difference between experimental and theoretical magnetic s u s c e p t i b i l i t i e s , the covalent character of the coordinate bond must also be considered. While covalency does not affect the spin, i t reduces any o r b i t a l angular momentum which might be present, and consequently, the effect of spin-orbit coupling. To measure the reduction i n the o r b i t a l angular momentum and the spin-orbit coupling, a parameter, K i s introduced. This has been variously c a l l e d the o r b i t a l reduction factor, the electron d e l o c a l i z a t i o n parameter, or the covalence factor. For complexes with A or E ground terms, the o r b i t a l reduction factor, K may be calculated from the expression: XA = XS0 2 v 10Dq (5.1) 10Dq 138 ,where = experimental atomic s u s c e p t i b i l i t y , X c n = spin-only value of the s u s c e p t i b i l i t y at the same temperature, X = spin-orbit coupling constant of the free ion, o 10Dq = d-orb i t a l energy difference i n cubic symmetry, 2 —1 2 N/3 = 0 . 2 6 0 7 3 cm . erg/gauss mole, a = I4 f o r E terms, a = 8 f o r A terms. The properties of the paramagnetic center are more conveniently expressed i n terms of Meff» ^ n e "effective magnetic moment", which i s related to the atomic s u s c e p t i b i l i t y by the equation: A 3 k T = — ™ ( 5 . 2 ) Substituting the values of the physical constants into t h i s equation and rearranging, we obtain the more f a m i l i a r equation: "eff = 2'828>fV ^ 4 . 5 . 2 Results and Discussion Copper(li) Complexes Magnetic s u s c e p t i b i l i t y data were obtained over the temperature range 320 - 90°K f o r the eopper(ll) tetrakis(pyridine) complexes, Cu(py)^X,> (X = S0 3F~, C10^", BF^~, CH^CgB^SO^) and f o r Ni(py)^(S0 3P) 2. The magnetic moments are given i n Tables XXXI, XXXIII which include the l i t e r a t u r e values of si m i l a r complexes. The experimental data on the complexes, giving the v a r i a t i o n , i n the magnetic s u s c e p t i b i l i t i e s as a function of temperature 139 are presented i n Table XXXH. A l l previously reported magnetic moments of the copper(ll) complexes were obtained at a single temperature only. Anion, X Magnetic Moment, Me£f Weiss Constant, 0 Reference CIO. BP, SO^F CH-C^H, S 0 -3 6 U 3 NO ~ CP 3COO CCl^COO ReO. " 1 . 8 2 1 . 8 0 1 . 8 8 1.81+ 1 . 7 8 1 . 8 2 1 . 7 2 1 . 8 2 1 . 8 8 1 . 8 7 1 . 8 5 1 . 8 2 1.91+ 1 . 8 0 0 ° ( 3 ) 3 ° ( 5 ) - 3 ° ( 3 ) - U ° ( 7 ) t h i s work 3k 35 t h i s work 31+ t h i s work 30 t h i s work 51 98 83 98 77 76 Note: Standard deviations of Weiss constants are given i n brackets. TableXXXE Room Temperature Magnetic Moments of Cu(py)^X 2 Complexes. 140 Table XXXII - Magnetic S u s c e p t i b i l i t y Data f o r Cu(py). X ? and Ni(py). (S0-.F) Temp (°K) 1 0 6 % (cgs) 1 0 6 % A (cgs) ^ f f 3 1 3 . 5 N 1 . 9 5 6 7 1 3 2 7 . 0 1 . 8 2 2 9 3 . 2 2 . 1 1 3 9 1 4 1 5 . 3 1 . 8 2 2 7 5 . 5 2 . 2 5 0 0 1 4 9 1 . 8 1.81 2 4 3 . 8 2.6311 1 7 0 5 . 9 1 . 8 2 2 1 2 . 0 3 . 0 5 4 5 1 9 4 3 . 7 1 . 8 2 1 8 3 . 8 3 . 5 2 6 3 2 2 0 8 . 7 1 . 8 0 1 6 4 . 0 4.0.737 2 5 1 6 . 3 1 . 8 2 1 4 5 . 3 4 . 5 5 4 5 2786. i 1 1 . 8 0 1 2 5 . 0 5 . 3 8 6 2 3 2 5 3 . 6 1 . 8 0 1 0 3 . 0 6 . 5 3 8 4 3 9 0 0 . 9 1 .79 9 1 . 1 7 .7541 4 5 8 3 . 9 1 . 8 3 diamagnetic correction = - 2 8 9 x 10 cgs. - 6 TIP correction = 6 1 . 3 x 10~ cgs. 3 0 9 . 2 1 .9721 1 3 5 6 . 7 1 . 8 3 2 9 2 . 0 2 . 0 8 5 1 1 4 2 2 . 1 1 . 8 2 2 7 2 . 5 2 . 2 7 6 2 1 5 3 2 . 7 1 . 8 3 2 4 4 . 9 2 . 5 5 4 2 1 6 9 3 . 7 1 . 8 2 2 1 4 . 6 2 . 9 7 1 2 1935 .1 1 . 8 2 1 8 5 . 7 3 . 4 7 5 2 2 2 2 6 . 9 1 . 8 2 1 5 7 . 2 4 . 1 9 6 2 2 6 4 4 . 3 1 . 8 2 1 3 2 . 4 5 . 0 4 7 6 3 1 3 7 . 2 1 . 8 2 9 8 . 4 7 . 0 8 9 3 4 3 1 9 . 3 1 . 8 4 diamagnetic correction = - 2 7 3 • 6 x 10~ cgs. - 6 TIP correction = 5 8 . 7 x 10~ cgs. 11+1 o 6 6 Temp ( E) 10 % ( CS S) 1 0 \ (°SS) /* 3 1 6 . 4 1 . 9 8 6 3 1 3 2 7 . 3 1 . 8 3 2 9 4 - 6 2 . 1 7 8 6 1 4 3 3 . 7 1 . 8 4 2 7 U . 3 2 . 3 4 5 2 1 5 2 5 . 9 1 . 8 3 2 4 4 . 0 2 . 7 2 3 2 1 7 3 5 . 2 1 . 8 4 2 1 5 . 6 3 . 1 9 0 9 1 9 9 8 . 6 1 . 8 6 1 8 U . 3 3 . 7 8 6 8 2 3 2 4 . 0 1 . 8 5 1 6 5 - 3 4 . 2 3 5 4 2 5 7 2 . 3 1 . 8 4 1 4 3 . 2 4 . 9 2 7 4 2 9 5 5 . 4 1 . 8 4 1 2 4 . 2 5 . 6 8 9 9 3 3 7 7 . 5 1 . 8 3 diamagnetic correction = - 2 8 6 . 8 x 10 cgs. TIP correction = 5 9 . 1 x 10 ^  cgs. Cu(py) l f(CH 3C 6H uS0 3) 2 3 1 2 . 1 1 .3521 1317.1 1.81 2 9 2 . 8 1 . 4 9 5 8 1 4 2 0 . 9 1 . 8 2 2 6 8 . 3 1 . 6 5 5 5 1 5 3 6 . 2 1 . 8 2 2 4 4 . 2 1 . 8 5 7 8 1 6 8 2 . 4 1.81 216.1 2 . 1 6 6 5 1 9 0 5 . 4 1.81 1 8 5 . 0 2 .5871 2 2 0 9 . 1 1 .81 - 6 diamagnetic correction = - 4 0 3 . 6 x 10 cgs. - 6 TIP correction = 6 3 . 2 x 10~ cgs. 142 Temp (°K) 1 0 6 % (cgs) 1 0 6 * . . (cgs) g A NiiE£)UCsp3P)2 3 1 2 . 7 6 . 7 6 1 3 3 9 1 8 . 8 3 . 1 3 2 9 2 . 8 7 . 2 5 3 7 4 2 0 1 . 1 3 . 1 4 2 7 3 . 6 7 . 7 6 1 2 4 4 9 2 . 0 3 . 1 4 2 4 3 . 8 8 . 7 0 9 3 5 0 3 5 . 4 3.1-3 2 1 5 . 6 9 . 8 2 4 5 5 6 7 4 - 7 3 . 1 3 1 8 5 . 2 1 1 . 3 2 8 6 6 5 3 6 . 9 3.11 1 6 5 . 3 1 2 . 6 9 9 2 7 3 2 2 . 6 3.11 1 4 4 . 7 1 4 . 4 5 8 7 8 3 3 1 . 2 3.11 1 2 5 . 3 1 6 . 6 0 1 0 9 5 5 9 . 2 3 . 1 0 1 0 7 . 7 1 9 . 8 8 9 0 1 1 4 4 4 . 0 3 . 1 4 diamagnetic correction = - 2 8 9 . 0 x 1 0 " cgs. - 6 TIP correction = 2 4 6 . 0 x 1 0 " cgs. 143 Since the room temperature magnetic moments of the complexes are a l l greater than the spin-only value, and l i e within the region 1.80 - 2.00, we can conclude that the complexes are almost c e r t a i n l y magnetically d i l u t e , and that there i s ne g l i g i b l e spin-spin coupling cha r a c t e r i s t i c of binulear 133 or polynuclear structures containing metal-metal bonds. For magnetically d i l u t e copper(ll) complexes with E ground terms, Curie law behaviour i s expected, i . e . , the magnetic moment should be indepen-dent of temperature. Such behaviour i s indeed observed f o r the perchlorate and tetrafluoroborate complexes, but, Cu(py)^(SO^F^ and Cu(py)^(CH3CgH^S03)2 show a s l i g h t decrease i n magnetic moment with decreas-ing temperature (Table XXXIl). A plot of 1/x^ versus temperature (Figure 25) reveals a l i n e a r relationship with a negative intercept, 0 . The Curie law therefore breaks down for these complexes and i s replaced by the Curie-Weiss law, which may be expressed by the formula: i n magnetically d i l u t e systems i f intramolecular antiferromagnetism i s responsible f o r the deviation from Curie law. However, many authors have used t h i s equation when Curie-Weiss law behaviour was observed under the vague assumption of antiferromagnetic interactions. Where there i s no evidence of such interactions, no pa r t i c u l a r significance can be assigned to the Weiss constant, which becomes simply an empirical quantity. (5A) where 9 i s the Weiss constant. Equation ( 5 « 4 ) may only be used to calculate the magnetic moment 1UU Curie-Weiss law behaviour may be observed f o r ions i n which the ground term i s degenerate or i s associated with a higher l e v e l s u f f i c i e n t l y close to be thermally populated. In such cases the resultant moment i s a combination of the moments associated with the in d i v i d u a l l e v e l s and i s therefore temperature dependent. Whatever the source of the deviation from i d e a l Curie law behaviour, the interactions responsible f o r i t are small, since the magnitude of a l l the Weiss constants are less than 5°K with a change i n Me££ no greater than 0 . 0 3 between 3 2 0 ° and 90°K. The magnetic moments of a l l the copper(ll) complexes studied were tested f o r f i e l d dependence to detect the presence of any ferromagnetic interactions. A l l r e s u l t s were negative. As the increase i n the magnetic moments above the spin-only value i s due primarily to the spin-orbit coupling, the magnitude of t h i s increase i s c l e a r l y dependent upon the symmetry of the excited states present and upon t h e i r energies, both of which depend on the stereochemistry present. However, the t o t a l contribution to the magnetic moment by t h i s mechanism i s so small, and the accuracy with which M e££ can he measured i s such that no stereo-chemical information can be obtained from the magnetic moment. Consequently, although the observed magnetic moments are consistent with a single unpaired electron, they cannot distinguish between a d ^ ^.2 ground state f o r an elongated tetragonal environment, or a d ground state f o r a compressed 2 tetragonal environment. N i c k e l ( l l ) Complexes The magnetic s u s c e p t i b i l i t y of Ni(py)^(S0 3F) 2 has been measured over the temperature range 313 - 107°K. The room temperature magnetic 11+5 moment i s compared with si m i l a r n i c k e l ( l l ) tetrakis(pyridine) complexes i n Table XXXIII. A l l published magnetic moments of the n i c k e l ( l l ) tetrakis(pyridine) complexes were obtained at a single temperature, except those of Belova, 136 Syrkin and Babaeva who report moments of the chloride, bromide, iodide, and thiocyanate complexes at 293°K and 78°K. The complexes reported at two temperatures show moments which are dependent on temperature. For example, at 293°K, Mgff f o r Ni(py)^Cl2 i s 3>^h> while at 78°K, i t decreases to 2 . 9 6 . This compares with a s l i g h t decrease observed here i n the moment of Ni(py)| +(S0 3F) 2 from 3-12+ at 293°K to 3.10 at 125°K (Table XXXIl). Since the n i c k e l ( l l ) ion possesses a A„ ground term, we would *-g i d e a l l y expect Curie law behaviour with moments independent of temperature. However, Ni(py)^(S0 3F) 2 shows Curie-Weiss law behaviour with a Weiss constant, 9 = -3°K (standard deviation = 1 ° ) . From the published molar s u s c e p t i b i l i -t i e s at 293°K and 78°K, the Weiss constant f o r the n i c k e l ( l l ) t e t r a k i s (pyridine) thiocyanate, chloride, bromide, and iodide were calculated to be - 1 1 u , - 8 U , - 5 , and - 2 UK. The cause of the deviation from Curie law behaviour i s as was described f o r the copper(ll) complexes. In a l l cases, the room temperature magnetic moments f a l l within the range 3«04 - 3 * 6 , which i s expected f o r octahedral spin-free n i c k e l ( l l ) , the increase over the spin-only moment of 2 . 8 3 depending on the magnitude of 138 the o r b i t a l contribution. 11*6 0 100 200 300 T E M P E R A T U R E ( ° K ) T E M P E R A T U R E ( ° K ) Figure 25 - l / x A versus Temperature for CuCpy)^ (C10 / [) 2 and Ni (py) 4 (S0 3F) 147 Anion, X Magnetic Moment, M e f f Reference SO^F" 3 . 1 4 t h i s work C 1 V 3 . 2 4 80 3 . 4 4 34 3 . 6 81 BF k" 3 . 0 4 80 3 . 4 0 34 CF^COO" 3 . 1 6 83 cci 3coo" 3 . 1 6 134 3.20 77 NCO" 3 . 0 4 80 3 . 1 7 135 NCS" 3 . 1 3 135 3.11 136 NCSe" 3 . 2 0 135 C l ~ 3-11 137 3 . 1 4 136 Br" 3 . 2 2 137 3 . 1 6 136 I " 3-21 137 3 . 1 5 136 Table IXXXIII- Room Temperature Magnetic Moments of Ni(p.y). X p Complexes. 11+8 U.6 ELECTRON PARAMAGNETIC RESONANCE SPECTROSCOPY The theory of electron paramagnetic resonance (epr) spectroscopy as i t relates to t r a n s i t i o n metal complexes has been well reviewed 1 39-1 lj.2 recently, and w i l l he repeated here only as i t i s applied to the copper(ll) systems studied i n t h i s work. 1+.6.1 Theoretical P r i n c i p l e s of Epr Spectra . N 9 The copper(II) ion with a d configuration, has an effective spin of S = -g- and associated spin angular momentum of m = +•§• leading to a doubly degenerate spin energy state i n the absence of a magnetic f i e l d . When a magnetic f i e l d i s applied, t h i s degeneracy i s removed. The low energy state has the spin magnetic moment aligned with the f i e l d and corres-ponds to the quantum number, m = --g-, while the high energy state, m = s s has i t s moment opposed to the f i e l d . A t r a n s i t i o n between the two different spin energy states occurs upon absorption of a quantum of radi a t i o n i n the microwave region. The energy, E, of the t r a n s i t i o n i s given by: E = he = g/3H where h i s Planck's constant, v the frequency of radiation, /3 the Bohr magneton, H the f i e l d strength, and g the Land'e s p l i t t i n g factor. The quantity g i s not a cscalar'-, but depends upon the orientation of the molecule containing the unpaired electron with respect to the magnetic f i e l d . I f the paramagnetic ion i s located i n a perfe c t l y cubic c r y s t a l s i t e , the g value i s independent of the orientation of the c r y s t a l and i s said to be i s o t r o p i c . In a c r y s t a l s i t e of lower symmetry, the g value i s orienta-149 t i o n dependent and i s said to be anisotropic. The z-direction i s defined as coincident with the highest-fold r o t a t i o n axis. In cubic symmetry, g = g = g , but when there i s an a x i a l symmetry, g 4 g - g and the & x & y z J J °z r °x & y usual notations are g = g , and g = g = g when the z-axis i s p a r a l l e l °z II x y x and perpendicular to the external magnetic f i e l d respectively. The unpaired d-electron i n an isolat e d copper(ll) ion possesses o r b i t a l angular momentum i n addition to the spin momentum, and these couple together to give a t o t a l angular momentum J . The Land! g factor i s given by: g - i J ( J + 1 ) + S ( S + 1 ) - L ( L + 1 ) 2 J ( J + 1) ( 6 . 2 ) where L and S are the quantum numbers f o r the t o t a l o r b i t a l and spin momenta of the ion. In an octahedral c r y s t a l f i e l d we expect a "quenching" of the angular momentum, and the g-value should be that of a single unpaired elec-tron, with g g = 2 . 0 0 2 3 1 9 . However, there w i l l always be some spin-orbit i n t e r a c t i o n which leads to a mixing of excited state configurations with the ground state configuration. I t i s t h i s contribution that makes the g factor a ch a r a c t e r i s t i c property of the ion, i t s oxidation state, and the molecular or e l e c t r i c f i e l d causing the "quenching" of the o r b i t a l motion. In symmetry, the g-values f o r the copper(ll) ion with a ^B-jg ground state are given by the expressions: % = e e [ 1 - k ^ \ ^ \ g — % g ) ] ( 6 - 3 ) 150 where K i s the o r b i t a l reduction factor and X q i s the spin-orbit coupling constant which f o r the free ion equals - 8 2 9 cm When an unpaired electron comes i n the v i c i n i t y of a nucleus with a spin, I, an int e r a c t i o n takes place which causes the absorption to be s p l i t into 2 1 + 1 components. The cause of t h i s s p l i t t i n g i s the nuclear spin-electron spin coupling a r i s i n g mainly from the Fermi contact term. Thus i n the case of copper(ll) which has I = 3 / 2 , the epr absorption spectrum consists of four l i n e s (Figure 2 6 ). The selection rules are Am = + 1 and A I L = 0 . m + 3 „ ^ V \ 1 \ " 3 / + 3/2 + 1/2 - 1/2 - 3/2 - 3/2 - 1/2 + 1/2 •+ 3/2 zero f i e l d magnetic f i e l d applied Figure 26 - Energy-level Diagram for S = -g- and I = 3 / 2 . 151 The hyperfine structure i n paramagnetic s a l t s can he interpreted using a simple Hamiltonian involving components of the electron and nuclear 1 JJ.3 spins. I f the anisotropy has a x i a l symmetry about the z-axis, then the spin Hamiltonian has the form: H = / 3 r g H S + g ( H S + H S ) l + A S I + B(S I + S I ) ( 6 . 5 ) H [_en z z 6i v x x y y'J z z v x x y y' ^ >/ where A and B are the hyperfine coupling constants p a r a l l e l and perpendicular to the applied magnetic f i e l d , respectively, i . e . , A = A|( = k^, B = A^ = A = A . The terms i n A, B correspond to the interaction between the nuclear x y magnetic moment and the magnetic f i e l d of the u n f i l l e d electron s h e l l . 1) i) i Bleaney has shown that including magnetic hyperfine interaction to second order, the resonance f i e l d i n a x i a l symmetry i s approximately given by: H = H - I o KmT B £p hg2P2n 2 ^ "o ,2 „ 2 V 2 2 g s i n 0 c o s 0 uij. ( 6 . 6 ) -p\ \ K / \ g j where H o = gelHr'e ; gQ = g-value of a free electron = 2 . 0 0 2 3 1 9 g T T T^2 2 = . 2 2 2- „ 2 2 . 2 f l H = ; K g = A g cos 0 + B g s i n 0 e e0 102 6 i s the angle between the molecular symmetry axis and the di r e c t i o n of the applied magnetic f i e l d . For 6 = 0 ° and/or 6 - 9 0 ° , the term i n 6 disappears. Substituting Yiv = ge/3Hg ; H Q = g e H e ; G = S G„ = S i ; G = £i_ and AH = H - H , g e S e Equation ( 6 . 6 ) becomes: \G / G l+HGV K 2 / For 0 = 0 ° , A H N * H e ( - - 1 r _ A - n - { i d + D - m,2} ( 6 . 8 ) V G / G|( 2HeG(| For 6 = 9 0 ° , 2 „2 \ G X / G X LH EG X 2 Putting D = \ ( 1 . _ -A ; E = ; F = ; f(lym) = I ( l + 1 ) - n^'2 " \ ' G n 2 H e G n 153 Equation (6.8) becomes: AH,, = D - EnLj. - P.f(l,m r) ( 6 . 1 0 ) Also, A H / ^ + A H / - 1 1 1 ^ = 2D - 2F.f(l,m T) ( 6 . 1 1 ) ( + m l ) . A H W = " 2 E m I ( 6 . 1 2 ) The three unknowns D. E. and P can be solved by substituting into equations ( 6 . 1 0 ) - ( 6 . 1 2 ) the values of A H f o r the different nuclear spin angular momentum quantum numbers. Thus, g^ and A can be determined d i r e c t l y from the epr spectra. Theoretically, the value of B can be obtained from the p a r a l l e l spectrum by solving the three equations, but since |B| « |A| , the value of B i s very sensitive to small changes i n AH ( | , hence unreliable values of B would be obtained. The values of hyperfine coupling constants obtained by t h i s procedure are i n units of magnetic f i e l d . They may be converted to frequency units by the use of the equation: 1 S Z A(cm ) = A(gauss) x — x 9-3481; x 10 ^ ( 6 . 1 3 ) S e The i s o t r o p i c parameters g Q and A q may be obtained from the epr spectrum of the complex i n solution at room temperature. The f i e l d positions 104 of the resonance l i n e s where m^  = --g- and ' = - 3/2 may he measured from the room temperature spectrum and substituted into the second order equation: 2 H = H Q - - ^o_ | l ( l + 1) - n^ 2} (6.11+) 2H o where H Q = "^"-i A f i r s t approximation of g Q i s obtained from an estimated value of H q from the i s o t r o p i c spectrum and substitution into g Q = he/H"o/3. An approximate value of A Q may be obtained from the difference TL^^ - H ^ j^. These f i r s t approximations of g Q and A Q may be substituted into equation (6.14), and the calculation reiterated with the measured values of H u n t i l the g Q and A q parameters converged. This usually requires 4-- E> cycles. Using these values f o r g(| , g Q, A, and A Q i t i s possible to obtain g^ and B from the equations: ?o - 3 ( g | | + 2 g J ; A o - 1(A + 2B) (6.15) 4.6.2 Results and Discussion The epr spectra of the series of Cu(py)^X 2 complexes (X = CIO^ BF^", SO^F", N03", CH^CgH^SO^", CF^COO") were obtained i n CH 2C1 2 + 10% pyridine solution at room temperature and at l i q u i d nitrogen temperature Figure 27 - EPR Spectrum of Cu(py), (j30_F)_ Solution at Room Temperature. 157 and t y p i c a l spectra obtained under these conditions are shown i n Figures 27 and 2 8 . The room temperature spectrum consists of four l i n e s due to the copper nuclear hyperfine interaction. The asymmetrical band shapes are a re s u l t of tumbling of the paramagnetic ions i n solution which tends to average out any anisotropies. In a p o l y c r y s t a l l i n e sample, or a glass, the molecular axes of symmetry are randomly dist r i b u t e d and the ove r a l l resonance l i n e shape i s the wisup1erp©sitionge of the contributions of the in d i v i d u a l resonance 1U5 centers. Sands has shown, that f o r such a system, the derivative spec-trum shows a set of weak l i n e s at the resonance f i e l d , H(| , corresponding to those molecules with t h e i r symmetry axes p a r a l l e l to the applied magnetic f i e l d , and a set of strong l i n e s at the resonance f i e l d , H x corresponding to those molecules with the symmetry axes perpendicular to the applied f i e l d . For a x i a l symmetry i n the Cu(yy)^X^ complexes, two l i n e s centered at g and g^, respectively, are expected. Each l i n e w i l l be s p l i t by the copper nuclear hyperfine in t e r a c t i o n into four components. This i s c l e a r l y observed i n Figure 28 where three of the four g|( components are v i s i b l e ; the fourth i s hidden by the g^ peaks which overlap i t . The s p l i t t i n g on the g^ peak i s not observed because of the superimposed nitrogen superhyper-fine s p l i t t i n g . Diagramatically, the spectrum of the frozen solution may be represented as Figure 2 9 . The superhyperfine structure (shfs) on the perpendicular components arises from the inte r a c t i o n of the nitrogen nu c l e i i n the pyridine molecules with the unpaired electron i n the copper ion. Since there are four equiva-lent nitrogen atoms surrounding the copper ion, each with spin = 1, we 158 A B 3/2 /2 y 2 Figure 29 - Line Diagram of the Epr Spectrum of Cu(py)^X 2 i n Frozen  Solution (without nitrogen superhyperfine s p l i t t i n g ) . expect nine shf l i n e s with a r e l a t i v e i n t e n s i t y r a t i o 1 : 1 + : 1 0 : 1 6 : 1 9 S 16 : 10 : h : 1 on each copper(ll) hyperfine component. Nine l i n e shfs i n the perpendicular spectrum have been observed previously"* ^  f o r many other copper(ll) complexes with four equivalent neighbouring nitrogen atoms and no apparent resolution of the perpendicular copper hyperfine components. In contrast to t h i s , about thirteen peaks were evident i n the perpendicular spectrum of a l l the copper(ll) complexes studied here. The excess of 2+ nitrogen shfs has been interpreted f o r Cu(py)^ complexes i n terms of 1 51 152 "extra absorptions". ' When the p a r a l l e l hyperfine coupling constant of copper i s much larger than the perpendicular hyperfine coupling constant, then, f o r certain values of g, A and m^,. there may be absorptions within the range 0 ° < 6 < 9 0 ° i n addition to those expected f o r 6=0° and 0 = 90°• The extra absorbances caused by such angular anomalies have been reported i n copper(II) acetylacetonates. 159 Schneider and Zelewsky also observed a large number of super-hyperfine l i n e s i n the single c r y s t a l spectra of Cu(py)^(lT0 3)2 and Cu(py)^(CH3CgH^S03)2 doped into the isomorphous platinum and cadmium s a l t s , respectively. They attributed t h i s to the presence of two magnetically non-equivalent complexes within the unit c e l l , where the angles between the tetragonal axes are 7 8 ° and 3 6 ° f o r the n i t r a t e and p - t o l y l s u l f a t e complexes, respectively. Such an explanation of the large number of super-hyperfine l i n e s does not apply to t h i s work however, since the molecular axes are randomly oriented i n a frozen solution. A plausible explanation of the 13 superhyperfine l i n e s observed i n the frozen solution "is that i t arises from an accidental overlap of the nitrogen shfs superimposed on each of the four resolved hyperfine i n t e r -action. I f there i s complete resolution of the nitrogen shfs superimposed on the copperthf s^ then there should be a t o t a l of 3& peaks i n the perpen-dicula r spectrum. Because of the close s i m i l a r i t y of the copper hyperfine and nitrogen superhyperfine s p l i t t i n g , there w i l l i n e v i t a b l y be some accidental equivalence of shf structure, thus decreasing the number of bands —Ii —1 resolved. A v a r i a t i o n i n the nitrogen shfs between ( 8 - l 5 ) x 1 0 c m within the perpendicular components of the copper hyperfine structure i n a t y p i c a l spectrum, supports t h i s p o s s i b i l i t y of accidental overlap of superhyperfine l i n e s . Ligand superhyperfine structure i s also observed on the low f i e l d + 3/2 -*—>• + 3/2 copper hyperfine component. In t h i s case, the expected nine nitrogen shf components in d i c a t i n g four equivalent nitrogen donor atoms about the copper ion, are present. The nitrogen superhyperfine coupling ii. 1 constant i s between (10 . 5 - 11-5) x 10~ 4 cm" i n a l l the complexes. 160 The epr parameters calculated from the spectra of the complexes as glasses and solutions using the second order equations are given i n Table XXXIV. x 10 cm Anion, X g|( g x g Q |A| |B| |AQ| C 1 0 ^ ~ 2 . 2 6 3 2.O4O 2 . 1 1 4 2O4 13 77 BF^~ 2 . 2 ^ 6 2 . 0 4 3 2 . 1 1 4 207 14 78 SO 3F" 2 . 2 5 3 2 . 0 5 4 2.121 209 5 73 N 0 3 ~ 2 . 2 5 8 2 . 0 6 1 2 . 1 2 7 209 4 * 72 CB^CgH^SO^ 2 . 2 6 0 2 . 0 6 2 2 . 1 2 8 207 6 * 65 CF,C00" 2 . 2 6 2 2 . 0 8 5 2 . 1 4 4 205 2 * 67 B has opposite sign to BF^ , C 1 0 ^ , S 0 3 F complexes. Table XXXIV - EPR Parameters of Cu(py)^X„ Complexes i n CH^Cl^ + 10% Pyridine Solution. The values f o r g and A were calculated from frozen solution spectra, g Q and A q from the room temperature solution spectra, and from these, g and B were obtained using equation ( 6 . 1 5 ) . 161 This procedure involves the assumption that the freezing of the solution does not change the l o c a l c r y s t a l f i e l d i n such a manner that would affect the i s o t r o p i c A q and g Q values. However, t h i s i s not s t r i c t l y true, as temperature dependence of the hyperfine s p l i t t i n g has been observed for 156 copper(ll) complexes i n solution. F a l l e and Luckhurst have suggested that t h i s may lead to an error i n the value of B as high as 20%. An important ramification of t h i s temperature dependence of hyperfine parameters i s that the commonly used assumption: A Q = A/3 + 2B/3 i s s t r i c t l y v a l i d only at the temperature at which a l l parameters are measured. Since i t i s often not possible to measure B and g^ from the spectrum of the frozen solution, i t i s s t i l l quite common practice to use equation (6.10) to calculate these parameters, ' ' ' ^ i n spite of the temperature dependence of the hyperfine parameters. I t should be noted that the absolute sign of A cannot be deter-mined experimentally, only the magnitude, but i t i s expected to be negative on theoretical grounds."^ 0 I f we r e l a t e the g-values i n Table XXXIV by the expression, G = g|( -2/g^ -2, then we obtain G values ranging from 6.6 to 4.2 f o r a l l the CuCpy^Xg complexes except f o r the trifluoroacetate complex where 121 133 G = 3«1« Hathaway and B i l l i n g ' have found that where the lowest G i s greater than 4.0, then the l o c a l tetragonal axes i n the copper(ll) complex are aligned p a r a l l e l or only s l i g h t l y misaligned, and i s consistent with an elongated tetragonal-octahedral stereochemistry f o r the copper(ll) ion 162 with a dx2_y2 ground state. This i s c l e a r l y the case f o r almost a l l the complexes studied here. I f G < 4 . 0 , s i g n i f i c a n t exchange coupling i s present, and the misalignment i s appreciable. In the only case where we observe G < l+.O, i . e . , f o r Cu(py)^(CF 3COO) 2 > the infrared spectrum of the complex i n CR" 2C1 2 + 1 0 % pyridine solution shows that the species i n solution i s predominantly Cu( p y ) 2 ( C F 3 C 0 0 ) 2 . Since the l a t t e r complex i s probably 82 dimeric with bridging trifluoroacetate groups, then i t i s not surprising that we observe G = 3«1 f o r the complex i n solution where exchange coupling should be present i n the dimeric species. The absence of exchange coupling i n the epr spectra of the fl u o r o s u l f a t e , perchlorate, tetrafluoroborate and p - t o l y l s u l f a t e complexes i s consistent with the magnetic s u s c e p t i b i l i t y data. The effect on copper(ll) epr parameters, of varying the a x i a l ligands has been studied quite extensively f o r copper(ll) bis(acetylacetonate) "1 "I ^ 8 and various substituted derivatives, , and copper(ll) b i s ( d i e t h y l d i t h i o -\ 159 carbamate;. The r e s u l t s of these studies show that as the a x i a l i n t e r -actions increase, g Q w i l l increase, and the magnitude of |AQI and iBl w i l l decrease. Thus, i f we arrange the g Q and IAQI values i n Table XXXIV i n increasing order of g Q, and decreasing order of IA Q| , we should obtain a series f o r increasing a x i a l i n t e r a c t i o n or coordinating strength: CIO, ~ BF, < SO^F < CH0C^H, S 0 _ 4 4 3 3 6 4 3 In spite of the p o s s i b i l i t y of a r e l a t i v e l y large error i n B due, to the temperature dependence of the hyperfine s p l i t t i n g , i t i s p a r t i c u l a r l y g r a t i f y i n g that a series obtained from consideration of B agrees with the 1 6 3 one derived from g Q and A. The v a r i a t i o n i n g^ for the diff e r e n t complexes also follows exactly the same trend as g Q. The trifluoroacetate and n i t r a t e complexes are not included i n the spectrochemical series derived above because of pyridine di s s o c i a t i o n to give the bispyridine complexes i n CHgClg + 1 0 % pyridine solution. This makes any comparison with the t e t r a k i s (pyridine) complexes, i n v a l i d . The use of epr spectroscopy and magnetic s u s c e p t i b i l i t y measure-ments on the series of Cu^py^Xg complexes provide two independent i n v e s t i -gations of the magnetic properties of the copper(ll) complexes. The 160 parameters obtained by these two techniques are related by the equations: u e f f = ^ 1 " 2«\/™Vz) (6.16) g = g e( 1 - 2K 2X Q/lODq) ( 6 . 1 7 ) M e f f = e oVs(S + 1) ( 6 . 1 8 ) Using equation ( 6 . 1 8 ) , and substituting S = for copper(ll) and the g Q values obtained from the room temperature epr spectra, we can compare calculated from epr data, and Mef£ experimentally obtained from bulk magnetic s u s c e p t i b i l i t y measurements. The resu l t s are summarized i n Table XXXV. The magnetic moments obtained by the two independent techniques agree within + 0 . 0 2 u n i t s , which i s within the experimental error i n the values. This agreement serves to confirm the v a l i d i t y of the conclusions 161+ Anion, X g Q M e f f ( c a l c . ) M e f f(expt.) CIO, ~ 2.111+ 1 . 8 3 1 . 8 2 BF, ~ 2.111+ 1 . 8 3 1.81+ S0 3F" 2 .121 1.81+ 1 . 8 2 CH3C6H^S03" 2 . 1 2 8 1.81+ 1 . 8 2 Table XXXV - Comparison of Magnetic Moments of Cufpy^X^ Complexes obtained from EPR and Magnetic S u s c e p t i b i l i t y Data. derived from the epr data. Since the epr data are obtained from solutions, and the magnetic s u s c e p t i b i l i t y data are obtained from the complexes i n the s o l i d state, we can reaffirm the assertion that there are l i t t l e changes i n the complexes when they are dissolved i n CHgClg + 1 0 % pyridine solution. Comparison of the energy separation between excited states obtained from electronic spectra (section 1+.1+.2) with the g Q values show that they follow the same trend. By using equation (6 . 17 )» i t i s then possible to calculate the o r b i t a l reduction f a c t o r ,K f o r the copper(li) complexes. The re s u l t s of t h i s c a l c u l a t i o n are given i n Table XXXVI, assuming 10Dq = electronic band maximum, v i n CH o01 o +10% pyridine max d. c. _ i solution, \ =.--829 cm , and g = 2 . 0 0 2 3 1 9 . 165 Anion, X v (kK) g maxv ' °o cio^~ 1 7 . 7 9 2 . 1 1 4 0 . 7 7 BF^~ 1 7 . 9 9 2 . 1 1 4 O.78 SO^F" 1 7 . 1 2 2.121 O.78 CB^CgH^SC^- 1 6 . 5 6 2 . 1 2 8 0 . 7 9 Table XXXVI - Orbita l Reduction Factor f o r Cu(py)^X,-, Complexes. The res u l t s show a rather constant value f o r the o r b i t a l reduction factor, K . This indicates that the v a r i a t i o n i n the g values i n these ' °o complexes i s dependent mainly on the extent of the d-orbital s p l i t t i n g , and that the difference between the Cu^py)^, complexes studied i s largely a c r y s t a l f i e l d effect rather than a covalent effect. In other words, the difference i n covalence i n the- metal-ligand coordination between the complexes studied, i s small. 166 U.7 SUMMARY AND CONCLUSIONS In t h i s work, we set out to prepare and characterize complexes of the f i r s t - r o w t r a n s i t i o n metal bi s f l u o r o s u l f a t e s . We have succeeded i n preparing the tetrakis(pyridine) complexes of n i c k e l ( l l ) , c o p p e r ( l l ) , and z i n c ( l r ) fluorosulfates. From infrared and electronic spectroscopy, and magnetic studies, we have shown conclusively that the n i c k e l ( i i ) and copper(ii) complexes have tetragonally distorted octahedral structures with four pyridine mole-cules i n a square plane, and the two fluorosulfate groups coordinated i n a trans configuration. The s i m i l a r i t y of the infrared spectrum of 2L: Zn(py)^(SCy) 2 with those of Ni(py)^(SC> 3F) 2 and Cu(py)^(S0 3F) 2 suggests a s i m i l a r structure f o r the z i n c ( l l ) complex. A further purpose of t h i s work was to estimate the strength of the fluorosulfate coordination to metals r e l a t i v e to other anions. In order to do t h i s i t was necessary to prepare and characterize related pyridine complexes. In so doing we have also learned much about the coordination strength of a number of anions. In p a r t i c u l a r , our r e s u l t s indicate that on the whole, the coordination strength of the fluorosulfate ion towards t r a n s i t i o n metals i s s i m i l a r , but s l i g h t l y stronger than that of the perchlorate and t e t r a -fluoroborate, but d e f i n i t e l y stronger than the very weakly basic hexafluoro-phosphate anion, and weaker than the perrhenate, p - t o l y l s u l f a t e , and trifluoroacetate anions. Evidence for t h i s comes from the r e s u l t s of the physical studies performed i n t h i s work. 167 The infrared mull spectra of the fluorosulfate complexes show the anions to he "semi-coordinated" i n a s i m i l a r manner to the perchlorate and tetrafluorohorate anions. However, the SO^F coordination appears to he s l i g h t l y stronger than ClO^ , or B F ^ coordination, since the anions are coordinated i n j^Zn(py)^(S0 3F) 2J, but not i n j^Zn(py) J ( C 1 0 ^ ) 2 and j^Zn(py) J ( B F ^ ) 2 ? ^ E l e c t r i c a l conductivity measurements on the Cu(py)^X 2 complexes i n a c e t o n i t r i l e solution gives the following ordering of anions: BF^ ~ ClO^" < FSO^" ( CH^CgH^SO^" < NO^ " < CF^COO" The molar conductances indicate the f i r s t three anions to be s i g n i f i c a n t l y weaker coordinated than the other three anions. 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