SPECTROSCOPIC STUDIES- OF LITHIUM ION SOLVATION IN NON-AQUEOUS SO'LVENTS Thesis for the Degree of Ph,’ D“ MICHIGAN STATE UNIVERSITY PAUL R. HANDY 1 9 72 Date 0-7639 This is to certify that the thesis entitled SPECTROSCOPIC STUDIES OF LITHIUM ION SOLVATION IN NON-AQUEOUS SOLVENTS presented by Paul R. Handy has been accepted towards fulfillment of the requirements for Ph . D . degree in Chemistry / / {@0/ Major professor 56; H7 ([77/ I I I I I I I ‘ __ ~‘ :3- BIN‘;ING BY "GAB & SBNS' * 800K NNDERYINC. LIBRARY BINDERS I .' srmsropt, m t; A{ ..,, “I: ~- I I cmsgfi L I B R A R Y Michigan State University seams-scene : Is The purpose c‘II red study of alkali to include several nique of litht‘m-7 | ties in varl ous n Far infrared and sodium salts l cf lithium salts ' 3.5-d1methy1pyr1d ¢ 5. the Ion-solver t 'n -1 03%cm for1 ammonium salts a fr‘eczuem:1«.=.g of t “tent by the “E C” the pyridine Indicate the f C Ehe 11 th 1 I. I "A 1th \ 8W 1‘50 US SOlVen Ta ~ “'5": °? 0.62 for T“! g‘l‘q ‘ 1.1m brOmir 1.2 h ' ~o .m do“? ABSTRACT SPECTROSCOPIC STUDIES OF LITHIUM ION SOLVATION IN NON-AQUEOUS SOLVENTS By Paul R. Handy The purpose of this study was to extend the far infra- red study of alkali metal and ammonium salts in pyridine to include several pyridine analogs and to apply the tech- nique of lithium-7 NMR to study lithium electrolyte solu- tions in various non-aqueous solvents. Far infrared spectra of a number of lithium, ammonium and sodium salts have been measured in h-methylpyridine and of lithium salts in pyridine, 3-methy1pyridine, 2-h, and 3.#-dimethylpyridine and 2-chloropyridine. The frequencies° of the ion-solvent vibration band were found at about 390 to 340 em"1 for the lithium salts, at 200 cm'1 for the ammonium salts and at 175 cm"1 for the sodium salts. The frequencies of the solvation bands are influenced to some extent by the nature and position of the substituent group(s) on the pyridine ring. In the case of halide salts, the data indicate the formation of contact ion pairs. The lithium-7 NMR chemical shifts at 59.6 MHz of lithium bromide and perchlorate solutions in several non- aqueous solvents have been determined over a concentration range of 0.02 to 0.5 molar and have been found to be linear for all solutions except for lithium bromide in acetonitrile. Lithium bromide solutions exhibit chemical shifts of 0.3 to 1.2 ppm downfield of the perchlorate solutions in pyridine, ' V “‘.\ ”V. ”Mr—J acetone. tetra‘eyd: mallet differenc: sulf oxide and ace {on in non-aqueos lith solvent don: shift. Itixed solve Smonde'wridi 15 Preferential' In the IESpecti Specific 1 in tetramethyl, broadening of signifies for lithiUm p: aetIVe cOmpou Paul a. Handy acetone, tetrahydrofuran and acetonitrile solutions with smaller differences found between the two salts in dimethyl- sulfoxide and acetic acid. The chemical shift of the lithium-7 ion in non-aqueous solvents does not show the same dependence with solvent donor ability as does the sodium-23 ion chemical shift. Mixed solvent studies in water-acetonitrile and dimethyl- sulfoxide-pyridine mixtures indicate that the lithium ion is preferentially solvated by water and dimethylsulfoxide in the respective mixtures. Specific lithium ion-solvent interactions were observed in tetramethylguanadine as indicated by the extensive line broadening of the lithium-7 resonance in this solvent. Significant lithium-7 chemical shifts were observed for lithium perchlorate in solutions of the physiologically active compound pentamethylenetetrazole in nitromethane. n1 " “(‘1’ CIF n"“ a” .uvst v VV‘ 1“ 1r SPECTROSCOPIC STUDIES OF LITHIUM ION SOLVATION IN NON-AQUEOUS SOLVENTS By ”(h . Paul R? Handy A THESIS Submitted to Michigan State University in partial fulfillment of the requirements for the degree of DOCTOR OF PHILOSOPHY Department of Chemistry 1972 r”: ?:I The author Hi Professor Alexandr: encouragement and lopreclati on bars of 'the group t‘ris work. ThankslI his Interest in tr: taDr. King K. 'u‘or. 53m! the early ; 1‘5 PSYChologi cal attribution of or ‘15 gift or ent‘. .. «E Special than} 4.33718 Burkhardt a! m“ 7‘. ”99p apprecp .0 a "I 103’s ’ Under-st it asLimited . ACKNOWEDGMENTS The author wishes to express his sincere gratitude to Professor Alexander I. Popov for his patient guidance. encouragement and support throughout this work. Appreciation is extended to those past and present mem- bers of "the group" who have contributed in various ways to this work. Thanks are extended to Dr. Frank M. D‘Itri for his interest in the instigation and completion of this work, to Dr. Ming K. Wong for numerous enlightening discussions during the early part of this work, to Dr. Ronald Erlich for his psychological counseling. to Mr. Gene Kales for his contribution of original art, and to Mr. Mark Greenberg for his gift of enthusiasm. Special thanks are given to two staff members. Mr. Wayne Burkhardt and Mr. Eric Roach without whose cooperation this work would have been much more difficult. Deep appreciation is extended to my wife, Debra. for her love, understanding and encouragement throughout the past year. It is to her and to our families that this work is dedicated. iv HCRCD'CTICII SPRARED H’""’ AUA\ 7‘" .nRLRED EXPEZ SCLVENTL. Pyrid: 4-nett Pyrid: Other TABLE OF CONTENTS Page INTRODUCTION 0 O O O O 0 O O O O O O O O O 0 1 INFRARED HISTORICAL . . . . . . . . . . . . u INFRARED EXPERIMENTAL PART . . . . . . . . . 11 SOLVENTS O O O O O O O O O O O O O O O O 11 Pyridine O O O O O O O O O O O O O O 11 4-m0thy1pyridine o o o o o o o o o o 11 Pyridine-d5. o o o o o o o o o o o o 11 Other substituted pyridines . . . . 11 Nitromethane . . . . . . . . . . . . 11 Benzene O O I O O O O O O O O O O O 1'2 SALTS O O O O O O O O C O O O O O O O O 12 Alkali metal salts . . . . . . . . . 12 Tetralkylammonium salts . . . . . . 12 Preparation of Salt Solutions . . . . . 12 Instrumental Measurements . . . . . . . 13 Near Infrared . . . . . . . . . . . 13 Far Infrared O O O O O O O O O O O O 13 Laser Raman O O O O O O O O O O O O 1“ RESULTS AND DISCUSSION . . . . . . . . . . . 15 4-methy1pyridine. O O O C O O O O O O O 15 Solvent Vibrational Bands . . . . . . . 22 Other Substituted Pyridines . . . . . . 28 Pyridine . O C O C O O C C O O O O O O 33 NMR HISTORICAL INTRODUCTION AND THEORETICAL . . . . . . 45 EVOLUTION or NMR TECHNIQUES . . . . . . . . 52 WATER PROTON CHEMICAL SHIFTS . . . . . 53 NMR Relaxation Studies In Aqueous Elec- trolyte Solutions . . . . . . . . 55 V a‘:‘.e of Contents ( :r. m "21w 3?: Alkali Metal 1. Electrolyt r a .. '1'". I ' '7: afmlxbn LALJ I Selvents . . Pyridine h-methylpy‘ Other Subs Acetronitr Acetone . Acetic Aci. Dimethylsu Nitrometha INSIBUKENTAT Ir Table of Contents (con't.) Page RELAXATION STUDIES WITH "OTHER NUCLEI" . . 58 Alkali Metal Ion Chemrcal Shifts in Electrolyte Solutions . . . . . . . . . 60 NMREXPERIMENTAL............... 65 $01vent8 . o . a o 0 . Pyridine . . . . . 4-methy1pyridin . Other Substituted Py Acetronitrile . . . . Acetone . . . . . . . Acetic Acid . . . . . Dimethylsulfoxide . . Nitromethane. . . . . Tetramethylguanadine . Lithium Salts . . . . . Tetrabutylammonium Salts ldi .....D'"..0 . ne . 0 ........U)... 0..0........ 0\ 0‘ O O\ \1 SAMPLE PREPARATION . . . . . . O\ \} INSTRUMENTATION . . . . . . . . NMR RESULTS AND DISCUSSION . . . . . . . . . . 70 CONCLUSION . . . . . . . . . . . . . . . . . . 102 LITERATURE CITED . . . . . . . . . . . . . . . 104 APPENDIX I . . . . . . . . . . . . . . . . . . 110 APPENDIX II . . . . . . . . . . . . . . . . . . 112 vi §‘. Ba'd inten ‘ solutions . Scectra 0“ ‘nutvlemrr on! methylryrifi' (a) Spectr‘. CEO-1% 032' molar lit‘ni :et'nylpyric‘ 0'1 m3! 3 f ‘ J oiution = I Lithium Ce w ~ .‘ 59'33'lpjri Figure 1. 10. ll. 12. LIST OF FIGURES Band intensities of lithium perchlorate solutions in 4-methylpyridine . . . . . Spectra of lithium perchlorate — tetra- butylammonium bromide solution in 4- methylpyridine. . . . . . . . . . . . . (a) Spectrum of 4-methylpyridine from 600-150 cm-l. (b) Spectrum of 0.31 molar lithium perchlorate in U- methylpyridine from 575-200 cm"1 0.1 mm, 3 H mylar beamsplitter. res- olution = # cm‘l . . . . . . . . . . . Lithium perchlorate spectra in a 3- methylpyridine-nitromethane mixture . . Spectra of some lithium salts in pyridine............... Spectra of 7LiCIOu and 6LiClOu in d5’pyr1d1ne . . . . . . o o . o . 0 o o Intensity of the #20 and #05 cm"1 bands in 1.96 E pyridine in benzene vs. LiClOu concentration . . . . . . . Lithium-7 chemical shifts of LiHr and LiClOu in various solvents . . . . . . Lithium-7 chemical shift versus solvent donor number . . . . . . . . . . . . . Lithium—7 chemical shift versus solvent Z value 0 O O O O O O O I O O O O O O O The change in the lithium-7 chemical shift of lithium perchlorate solutions in pyridine with added tetrabutzl- ammonium bromide . . . . . . . . . . . Lithium-7 chemical shifts in water- acetonitrile mixtures at 60.2 MHz . . . vii Page 19 21 2a 31 37 40 11,3 79 85 88 91 93 ":z cf Figures (C Fi-ixe nu .J. . .‘c Lithium-7 ' sulfoxide - Litmus-7 perchlorat trile vers The lithit LiClCu in tetrazole FION char Flow char List of Figures (con't.) Figure Page 13. Lithium-7 chemical shifts in dimethyl- sulfoxide-pyridine:mixtures. . . . . . . 95 14. Lithium-7 chemical shift of lithium perchlorate and bromide in acetoni- trile versus temperature . . . . . . . . 98 15. The lithium-7 chemical shift of 0-1 N LiClO in nitromethane-pentamethyl- tetragole solutions at 23.3 MHz . . . . 100 Al. Flow chart of program smooth . . . . . . 119 A2. Flow chart of subroutine start . . . . . 120 viii . Solvation 1: solvents (c Solvat i on i i e-Eet'nvl Ev: Solvent vi' changes To u-methyl p3; methane so Ll thium s 0 other S‘J’f‘g Pyridine“ Stituti on 'r . ‘ ‘1 relates for Vario‘. T '“1 till Urns; Chem Cal no 3‘73an in .. ling . Solutioah 71 ‘ ” Svihe c Table 1.. 6. 7. 10. 11. A1. LIST OF TABLES Page Solvation band positions in various solvents (cm'l) . . . . . . . . . . . . 5 Solvation band maximum frequencies in u'methyl pyridine 0 . 0 0 0 . 0 0 . 0 . 16 Solvent vibrational frequency (cm'l) changes for lithium salt solutions in 4-methy1pyridine and pyridine-nitro- methane solutions . . . . . . . . . . . 26 Lithium solvation band positions for other substituted pyridines . . . . . . 29 Pyridine--lithium ion band mass sub- stitution shifts (cm‘l) . . . . . . . . 33 T1 related processes. . . . . . . . . . 50 Observed lithium-7 chemical shifts of aqueous salts . . . . . . . . . . . . . 71 Diamagnetic susceptibility corrections for various solvents . . . . . . . . . 73 Lithium-7 chemical shifts in various solvents at 59.62 MHz vs. 9! LiClOu . . 7h Chemical shifts for lithium salt solu- tions in several solvents at 23.3 MHz 0 I O O O O O O O O O I 0 O 0 0 0 0 82 Temperature variation of the lithium-7 NMR linewidth in 0.5 fl LiClOu TMG solution . . . . . . . . . . . . . . . 101 Listing of program smooth . . . . . . . 121 ix .1.‘ ..-0 :11. ADJ“ *1 the. c ‘7te solutions have :me of the spec‘. :eiasyet. It is :r::ertie . such a? e‘:ilitv. as b unsure of the \ szlutions. These s 121:. solvent sepa: u) . 'I‘uu ct ion pairs 1 Classical tec tear. usually d .u. , '0 INTRODUCTION Although the chemical and physical properties of electro- lyte solutions have been studied for many years, the complex nature of the species in solution has not been fully elucida- ted as yet. It is known that the solvent physio-chemical properties. such as dielectric constant, dipole moment and sol- vating ability, as well as the salt characteristics. determine the nature of the various species found in electrolyte solutions. These solution species include free solvated ions, solvent separated ion pairs. solvated and unsolvated contact ion pairs and higher ionic aggregates. Classical techniques such as conductance and ion trans- port can usually distinguish between charged and uncharged species while colligative property studies give the total number of particles in solution which lead to average solute formula weights. All of these techniques measure bulk solu- tion properties and give little information about the chem- ical nature of species present in solution. For instance, it is nearly impossible to distinguish between contact and solvent separatedion pairs by the above techniques. Like- wise, attempts to study the extent of solvation by the deter- mination of the number of solvent molecules associated with a particular ion (the solvation number) usually result in ambiguous results. Tn recent yea fcrthe study 0f 1 :aticular. far 1! :zzniu: Salt 501‘ b . .2212 vibration : sclvation shell) . reieoendent nai: sailer degree 0.”. .. .A p frequenci e s a i .3. 8676? al cases .ectric constant .Nn. a3: hift to lower :t‘rter ions. It We? -.rates into t mated contact Recent work ..- is a very ser 3:";1‘2 411 ion in sol 139d ‘ .o correlate I.‘ i 5 ~g“eh‘ ..‘..S as well \.‘ obvious the . Lt «.“ug Q‘? and PP yu 3932M. "G. ”M “« Q {a "Q 2 In recent years the applications of spectroscopic techniques such as near and far infrared, Baman, electron paramagnetic resonance as well as proton, wideline and pulse nuclear magnetic resonance (NMR) have opened new approaches for the study of the species present in electrolyte solu- tions as well as the equilibria between these species. In particular, far infrared spectra of nonaqueous alkali and ammonium salt solutions reveal bands which appear to be due to the vibration of the cation in a solvent cage (inner solvation shell). The frequencies of these solvation bands are dependent mainly on the mass of the cation and to a smaller degree on the nature of the solvent. In most cases the frequencies are independent of the counterion. However. in several cases where solvents with moderate and low die- lectric constant were used, the solvation band was observed to shift to lower frequency when halide anions were used as counter ions. It was assumed that in these cases the anion penetrates into the cation inner solvation shell to form a solvated contact ion pair. Recent work in this laboratory has shown that sodium-23 NMR is a very sensitive probe of the environment of the sodium ion in solution. Sodium-23 chemical shifts have been used to correlate relative solvating abilities of a number of solvents as well as to study contact ion pair equilibria. It is obvious that the study of other ionic nuclei such as lithium-7 and potassium-39 would also lead to interesting results. This thesis describes two approaches to the study of nonaqueous electrolyte solutions. Infrared and Raman '=-L".‘.ques are app "' ‘I\‘b :nzzzne solutions gym-7 3.93 is i ar.‘ 13: pair fOI‘TTB Tanzus solvents . 3 techniques are applied to the study of pyridine and substituted pyridine solutions of alkali and ammonium ion salts. Also lithium-7 NMR is investigated as a probe into the solvation and ion pair formation properties of the lithium ion in 'various solvents. In 1965 Evans stuiies Of' “walk: cl. The? obserV! s;-e::ral region "hf as o.’ either the as dependent on t .‘Ltration was ascr viration and cons tztration in solut Etta-fled by Klan": ifilmonium (33‘: t'tlz-rof‘orr. and eye “Stands in the 7 3938 assigned t0 1, 4 ‘Jon an} . U; .. v H... arise 4'- A, v. t‘ .19 " .4‘eA f ‘ or the .“'7 a v 0; 801'? er :1: t. . v Nata a INFRARED HISTORICAL In 1965 Evans and Lo1 described their far infrared studies of tetralkylammonium halide salts in benzene solu- tion. They observed an absorption band in the 100 cm'1 spectral region which was not attributable to a vibrational mode of either the solvent or the solute. The band position was dependent on the mass of both the cation and anion. The vibration was ascribed to the direct cation-anion ion pair vibration and constituted the first report of an ionic vibration in solution. This work has only recently been extended by Klundt, _e__t g_1_.2 to the study of tertiary- alkylammonium (R3NH+) halide salts in carbon tetrachloride, chloroform and cyclohexane solutions. These workers observed two bands in the 72 to 111. and 132 to 198 cm-1 ranges, which were assigned to the hydrogen bond bending mode and the cation-anion stretching vibrations respectively. In 1966. Edgell, gt 31.3 reported far infrared bands of sodium salt solutions in tetrahydrofuran. These bands did not arise from either the salt or the solvent but appeared to be a function of the solvated cation itself. Bands were observed for the alkali metal and ammonium salts in a variety of solvents. A list of the solvents studied and the band frequencies observed is given in Table l. n F) (I) O F‘ <3 W V (1’ H O ’3 ! l3. invests .flotarn.a A.“ “LAC amid V ‘9. a: Table 1. Solvation band positions in various solvents (cm‘ ). Solvents Li+ NH4+ Na+ K+ Ref. onsoa #29 21L» 200 154 5 DMSO-d6 425 4 (Pr)280 #20 223 218 5 (Bu)280 425 225 220 5 2-pyrrolidone 400 218 206 145 6 l-Me-Z-pyrrolidoneb 398 207 205 140 6 acetone 425-409 212 195 7 acetone-d6 300-372 190 7 acetic acid 390 8 THF 413-375 192-184 142 3,4 Propylene c 401-384 184 185 141 9 carbonate Pyridine no 199 180 10 8- bands observed at 125 and 110 cm‘”1 for Rb+ and Cs+ resp. b- band observed at 106 cm"1 for Rb+. 0. bands observed at 115 and 112 cm"1 for Rb+ and 03+. ’"rese bandS. W 1:27-33 SOT-fem: delfi‘ T'" half-half”: t! "01' The SClvat‘. :en'al accuracy 0? 2:.iicated in all c In anion dependem :fhigh polarity a :2 near strength "z: ‘- .:_.sr ex.ent ace 2"!“ ‘-h~"0 . dependence azii is an excep‘ sting solvent 3‘ Since the l 3::"H» ““3313. Tsatc 311:: t ‘. d 1.2-Rea ~ ‘ Cara. “R 31 “\‘..5“ "\Oh u 0? t; These bands. which have become known as solvation bands. are characterized by their strong cation dependence and much weaker solvent dependence. The bands are intense, but broad, with half-height band widths equal to, or greater than 50 cm’l. The solvation bands follow Beer's law within experi- mental accuracy of the measurements. Isotopic substitution indicated in all cases that the band frequency varies inver- sely with the change in mass of the cation or the solvent. No anion dependence of band position is observed for solvents of high polarity and donor ability. Solvents of low polarity 3'4 and to a or donor strength, most notably tetrahydrofuran lesser extent acetone7 and propylene carbonate9 do show some anion dependence of the solvation band frequency. Acetic acid is an exception in that this low polarity, weakly do- nating solvent shows no solvation band anion dependence.8 Since the last review of spectroscopic studies11 in solutions, Tsatsas and Risen12 have reported two concentra- tion dependent far infrared bands at 195 and 160 cm'1 for sodium tetrabutylaluminate in cyclohexane as well as a Raman band at 202 cm"1 also attributed to the sodium ionic motion. These authors also observed far infrared bands for lithium, sodium and calcium ions in carboxylate- containing polymers.13 Dimethylformamide (DMF) solutions of lithium salts were studied by Lassigne and Blainelu by infrared and proton nuclear magnetic resonance (NMR). Their infrared studies showed bands at 420 and 365 cm‘1 which were attributed to solvation of the lithium ion by DMF. Splitting of the carbonyl band in these same solutions indicated that the ion-solvent maxim: occurs t :22: for lithium w fzrbzt'n lithium per 31.21123. 1 \n Day. at all. ‘31.: and t trahydro Z-i-C vibrational t aiiiticn of sodium :zlatir. of soditu 27:.cnerane mixturt si :slecular 01-h: Fi‘“n '_0ou. inn .0. 5 "1th inf “~11 a: 8 Etonel sols ""37 ”“8 perrom ”‘Hn ..... “3 ion CO mp1 0!..- 2“,flhb hob‘mtetxil- 1.qu “HE‘S in aceto. t Elissa. .0 ye the Wgrk Of Pau] h . 31 Sui. ‘\?Q‘ 7 interaction occurs through the oxygen atom. The solvation number for lithium with DMF was determined by NMR to be 4.3 for both lithium perchlorate and iodide in DMF-dioxane solutions. Day, g£_al,15 reported an interaction between sodium ions and tetrahydrofuran based on the observation that the C-O-C vibrational bands of tetrahydrofuran are split by the addition of sodium salts. These authors also studied the :xflvation of sodium tetrabutylaluminate in tetrahydrofuran- cyclohexane mixtures in which the solvation number of sodium was shown to vary with salt concentration.16 Kecki. 23 31- have applied normal coordinate analyses and molecular orbital calculations to solvation studies. Beginning with infrared and Raman studies of acetonitrile17 and acetone18 solvent vibrations in lithium salt solutions. they have performed normal coordinate analyses of the solvent- lithium ion complex. They have been able to predict the solvent-metal ion stretching band frequencies for several cations in acetone and acetonitrile. In order to further elucidate the changes in the solvent vibrational spectrum upon complexation, they successfully applied CNDO (complete neglect of differential overlap) calculations to eXplain the decrease in the acetone C=O stretch frequency and its increase in intensity upon complexation.19 Other near-infrared studies of ionic solutions include the work of Paul, 33 al.20 who studied the frequency shifts of dimethyl sulfoxide S=O and C-8 stretches in solutions of silver and lithium salts. They confirmed the solvation {Liar 03' 2 found i:eth;.'lsu1f‘oxide . gees-a of aqueous ate: and determin inthe salt hydra Ecuviere, it: 1: szlutions of al izissolvate by :i:ra:e"r.ane mixtu :$119 by fickirme Another appr: sixticns is the 5 32:26.21, a £2.21. il'i::yanate (113301 3.331111% on the tree components c 21 for lithium salts in number of 2 found by Maxey and Popov dimethylsulfoxide. McCabe and Fisher22 took near infrared spectra of aqueous alkali halide solutions referenced against water and determined a number of solvation parameters includ- ing the salt hydration volume and solvation number. Rouviere, gt al.23 studied pyridine vibrational changes in solutions of alkali metal salts. They found that lithium ion is solvated by h pyridine molecules in pyridine- nitromethane mixtures and confirmed the #20 cm'1 band reported earlier by McKinney and Popov.10 Another approach to the infrared study of electrolyte solutions is the study of polyatomic anion vibrations. Chabanel. §_t_ 2})?” studied the Cam vibration of lithium thiocyanate dissolved in various solvents. They found that, depending on the solvent, this band splits into at least three components corresponding to lithium thiocyanate ion aggregates. ion pairs or free solvated ions. The band com- ponents assigned to the ion aggregates and ion pairs were observed in dimethylcarbonate, while in dimethylsulfoxide the only band components present were those assigned to the free solvated ions and ion pairs. The authors found in- creased dissociation in the solvent order, dimethylcarbonate < ethylacetate < propylene carbonate < dimethylacetamide < dimethylsulfoxide. Edgell, 23; 11,25 studied the effects of different solvents on the vibrational spectrum of the tetracarbonylcobaltate anion in sodium tetracarbonylcobaltate solutions. The 1890 cm'l, C-O stretch band was observed to be quite symmetrical v. “V." flzethyl 5-5.1 f .13. 1:11 sates a Sc 2:72 and below the 21:: increases in 2271:: indicates ::ei'.'-.on resultin: :fthe scdiu'n ion b :h S .m soiuti ons re s , .t- u I; ll .32 crn‘l band I M‘- .:..it‘r.:um salt 5: .esscmplex. Th1“— 133 that the li thi mi] 10'] pairs .l‘ent ~¢;n11 and I y ' f. 9 in DMF, dimethylsulfoxide and in 5 per cent water-tetrahydrofuran, which indicates a symmetrical environment about the anion. However, in dimethoxyethane, pyridine, tetrahydrofuran and piperidine, the band is split with band components appearing above and below the 1890 cm"1 band. The magnitude of pertur- bation increases in the solvent order given above. This behavior indicates that there is increased asymmetry about the anion resulting from ion pairing effects. Replacement of the sodium ion by potassium and lithium in tetrahydro- furan solutions resulted in the same order of magnitude of the 1890 cm"’1 band complexity for the potassium salt, whereas the lithium salt solution spectrum in this region was much less complex. This latter case was interpreted as indicat- ing that the lithium tetracarbonylcobaltate is less exten- sively ion paired than the sodium or potassium salt in this solvent. Edgell and Lyford26 further observed that the complexity of the 1890 cm"1 tetracarbonylcobaltate band in tetrahydro- ‘ furan decreases with decreasing temperatures. At low tempera- tures it was possible to resolve the band into components corresponding to two kinds of asymmetrical anion environments which were assigned to contact and solvent-separated ion pair formation. The authors further concluded that both types of ion pairs contribute to the sodium solvation band. Pyridine and its analogs form interesting systems for the study of ionic salvation and complexation by spectro- scopic techniques. Infrared spectroscopic studies of alkali salt solutions in pyridine were reported by McKinney.27 .u Ar C'f v A v‘ ' U i O. '...-5 e .e- q . 1 u c l E alums -017 .~_._..:ss: M. S W‘ ..‘--A'l- -s -“.’. ‘c‘ ‘ Ln-vS: eaten é “TV-“4 or ‘i “-1 VAVH ‘="“-‘ 1‘ “yr .-‘w..g u “ ”J 10 The first part of his thesis was concerned with the deter- mination of the pyridine-iodine charge transfer complex formation constants for a series of pyridine analogs in various solvents. He was able to relate the formation constants directly to the pyridine basicities. Steric deviations were only observed for 2,6-dimethylpyridine. Infrared and Raman studies of pyridine and its mono- substituted and dimethyl analogs include complete fundamental vibrational analyses.28'29 Infrared studies also include a report on pyridine fundamental vibration spectral changes30 in solutions with hydrogen bonding solvents as well as far infrared studies of transition metal complexes with pyridine and some monosubstituted analogs.31 With the above information concerning donor properties, steric hindrance effects, and fundamental vibrational fre- quency dependencies of the pyridines. it should be possible to interrelate any changes in the solvation band behavior with changes in solvent properties. H EtriiirtE: ‘—-—— refined over 1: glass hell ci es p ...:I‘. I- L $39 .hvlpv "Q‘ Wu '?‘{J‘ "'““;r.° ~S Used ‘ u INFRARED EXPERIMENTAL PART SOLVENTS Pyridine: Fisher "certified" reagent pyridine was refluxed over barium oxide and~distilled through a 1 meter. glass helicies packed column. The purified material was stored in an amber bottle over barium oxide or molecular sieves (Fisher type 4A). The water content, as determined by Karl Fisher titration32, was found to be about 5 milli-~ molar. h-methylpyridine: The commercial product was obtained from Aldrich Chemical Company. Its purification and storage were analogous to that of pyridine. Karl Fisher titration showed that the water content was about 5 millimolar. Pyridine-d5: The NMR reagent grade material was obtained from Diaprep Inc. and was used without further purification. Other substituted pyridines: The other substituted pyridines used in this investigation were obtained from the Aldrich Chemical Company. They were purified by refluxing them over barium oxide for at least two hours. followed by fractional distillation. The distillation products were stored over barium oxide. Nitromethanes Matheson, Coleman and Bell. practical grade nitromethane was percolated through a 30 cm column of Dowex sow-X9 cation exchange resin and dried over Drierite ll :‘2:se'=-.-'eral hours “ashDrierite th: Trevater concent Tfish-31' titration. Zane-ov- D . U-uv. l'i‘ 'u'q. \ . . us..-"- 5 f \- ( were obtained a: \ 22". f‘thEr Fur: I I...‘ ‘Q rid SUeci ) F a:“ -ier ' S o< ~'. 12 for several hours. The solvent was then distilled from fresh Drierite through a 50 cm Vigreux column. The fraction boiling at 100.2°C was taken and stored in an amber bottle. The water concentration was found to be 0.017 molar by Karl Fisher titration. Benzene: Matheson, Coleman and Bell, ”chromatoquality", reagent grade benzene was dried over barium oxide before USE. we Alkali metal salts: Most of the alkali metal salts were obtained as reagent grade chemicals and were used with— out further purification, except for drying at 180 to ZOd’C for a minimum of forty-eight hours. Lithium iodide and thiocyanate were unstable at elevated temperatures and required special purification procedures which were reported earlier. Sodium and ammonium thiocyanate: and sodium tetra— phenylborate were dried under vacuum at 60°C for forty-eight hours. The preparation of lithium-6 salts has been described previously.27 Tetralkylammonium salts: Tetrabutylammonium perchlorate and bromide salts were obtained from Eastman Kodak and were dried at 60°C under vacuum prior to use. Preparation of Salt Solutions Because of the hygroscopic nature of the salts and solutions used in this study. care was taken to minimize exposure to the air. Weighings were conducted as quickly as possible with solvent and solution transfers being performed 381.131 0f t}, I is “MM . VJ“ «Ler tranh 2“ t tr‘ve ;: §-o . I‘OmetEr 15 lead 2:55 and the flask .‘zesoluticns were Lained at room temp Izstr'mental Pieasur riear Infrared: , -l .m spectral 2 '3'3‘1 225 Spectror-‘v we” POtassium :ell holder manufq ‘5‘1- \ c patolenzftn we. -;c1 o. ‘ 13 with pipets or syringes. For the preparation of most solu- tions, the salt was directly weighed into the volumetric flask and the flask was filled to the mark with solvent. The solutions were all prepared and their spectra were ob- tained at room temperature. Instrumental Measurements Near Infrared: The near infrared spectra in the #000 to 600 cm'1 spectral region were obtained on a Perkin Elmer Model 225 Spectrophotometer. The solution samples were held between potassium bromide salt flats in a standard demountable cell holder manufactured by Barnes Engineering Company. The cell pathlength was varied from 0.1 to 0.015 mm by the use of teflon spacers between the salt flats. Ear_Infrared: Spectral measurements below 600 cm"1 were obtained with either a Perkin Elmer Model 301 spectrophoto- meter or a Digilab FTS-l6 spectrometer. The characteristics and operation of the 301 spectrophotometer are given by Maxey.33 The FTS-l6 spectrometer is based upon a rapid- scan Michelson interferometer operated under computer control. The output of the interferometer, called an interferogram. is a sum of the Fourier components of the frequency spectrum. The Fourier transform is performed by a Data General, Nova computer which contains 12K of 16 bit core memory. Beyond interferometer control and spectrum computation, the computer is used to improve the spectrum signal-to-noise ratio through signal averaging of the interferogram. The computer allows the spectral data to be displayed as single beam ,_n| ‘ Ltd" V by n F. Tithe-e" u.” the instrument see zylar tears?”L e§.midioovers :15? and 230 to E :rofthe instru: tieentire range be zybe realized bei 11;?! energy radi; :fneinstrurent We manufacture Yost of the 3 713:5 of either 2 13:2 crn'1 I‘ESpec £395. With nominal E"“iiifd demounta -35. P‘lef‘n ul,d r» «Elly Here 813 tite' . A} i. i“ U 31Vent ad ‘35 . I . SOI‘jj hu I “5 were 1‘ e- W and the ‘. Ht‘éf - ' rtelativ: ifgas 3“ v ~aser N l y . W' T.“ ‘ v NS 9'... w “a“ '. 511.. .is. a 14 emission, transmission, absorbance, or log absorbance versus wavelength. The instrument covers the 600 to 50 cm‘1 region with three mylar beamsplitters of 3, 6, and 12 microns in thick- ness, which covers the nominal ranges of 600 to 150, 425 to 100 and 200 to 50 cm"1 respectively. The nominal resolu- tion of the instrument can be varied from 16 to 1 cm"1 for the entire range below 600 cm‘l. Half wavenumber resolution may be realized below 490 cm‘1 if care is taken to exclude higher energy radiation. Details of the theory and operation of the instrument may be found in the various manuals supplied by the manufacturer. Most of the spectra were obtained at nominal resolu- tions of either 2 or h cm’l, which gives a data point every 1 or 2 cm"1 respectively. Two types of sample cells were used with nominal pathlengths of 0.1 and 0.2 mm. These standard demountable cells were used with 2 mm polyethylene discs. Molded polyethylene cells from Barnes Engineering Company were also used, but were found to be inferior to the demountable cells for intensity measurements. Problems of solvent adsorption on the polyethylene windows and cell matching were reduced by taking the ratio of the solution spectrum and the previously stored spectrum of the cell itself. Relative band intensities were obtained from band areas as determined with a planimeter. Laser Raman: The Raman spectra were taken with a Spectra-Physics Model 700 Raman Spectrometer equipped with a U0 milliwatt, 6328 K helium-neon laser. .—. u- a.“ v:'}~o‘ ""$* A 1“ o ..-"..n‘.ufi;‘4 now 'i (I) x” y; 0 H. H 3 ‘f (h 2 inzeztse band 3' .5. d‘ti a weak b. 2:13 4210- oil 18 asxell as for rat Janis have been c Prd‘ " . A ‘N nine thylsulfc . .u‘ hag s. 8'11“ ‘zfi ‘ \gt ‘6 CI? +Y Una RESULTS AND DISCUSSION 4-methylpyridine Two basic criteria for solution spectroscopic studies are the absence of interfering solvent absorbtion bands in the region of interest and adequate solute solubility. The infrared spectrum of 4-methylpyridine below 600 cm"1 has an intense band at 490 cm'l, a medium intensity band at 512 cm"1 and a weak band near 210 cm‘l. The Lack of solvent bands near #00 cm"1 is unique for the pyridines used in this study as well as for most of the other solvents in which solvation bands have been observed. Pyridine, acetone, acetic acid and dimethylsulfoxide, all have solvent vibrational bands within 20 to 30 om-l of the observed lithium solvation band. This large spectral window combined with adequate salt sol- ubilities make h-methylpyridine a nearly ideal low polarity solvent for these studies. Solvent band frequencies observed in h-methylpyridine are presented in Table 2. In this solvent the lithium sol- vation band is observed at 390::3 cm"1 for salts of poly- atomic anions and at distinctly lower frequencies with halide anions. Replacement of lithium-7 by lithium-6 salts shifts the solvation band to higher frequency by approximately 25 cm'l. The shift in the solvation band frequency with the mass of the cation confirms that the cation is the major 15 16 Table 2. Solvation band maximum frequencies in 4—methyl- pyridine. LiCioLL 39li3cm’1 LiNo3 389 LiSCN 3H7 Liar4 387 Lil 384 LiBr 383 LiCl 378 6LiClO4 412 6LiBr 407 6LiCl hon usualou 200suom'1 NH43CN 207 NHQI 198 manna“ 198 NDQI 193 NaSCN 180:50m’1 NaClOu 180 NaBF4 176 NaI 17a .‘p‘i‘fiq VI ~.fo_. ‘m’fifi .' en— I...‘ .‘.i ‘ '- ..-l "‘1' B 11trMJU ‘ h .-":‘l 314:1 O“ . . 4 saw! ..-a A: “1"..flrll I. n n 13;" V. “ lptn two bands W1 "Inna : .1 . _.._L‘--: tEly eq ‘4 | ‘ a. . n ‘3. '2: :een reeOT W- I‘ ‘ . in“ “V’VWO .-¢‘-l-‘-I--. “‘ ' 3:: tr edl‘il‘. 1 0‘1 ‘)0 zen: with the ea Solvation k function: 2.7.1-...ecarbo“ 33"8" on barla ‘ Vt -~ tn Q v, ~ 11 fi‘ Phi -‘ “N{ 4- Q fly. 1 ‘ 011 o ~4. .:‘ «“1. m. 1? contributor to the observed vibration. A simple, diatomic Hooke's law calculation predicts a frequency shift of 30 cm"1 for a lithium-ion--b-methylpyridine species on lithium-6 substitution. The ammonium and sodium bands generally occur in the 200 and 180 cm'1 frequency regions respectively. In the case of ammonium thiocyanate, however, the solvation band is much broader than the corresponding bands for other ammonium salts. It appears that this band might be resolved into two bands with maxima at about 220 and 200 cm'1 of approximately equal intensity. A similar complication has been reported earlier for this salt in propylenecarbonate solution.9 The solvation band frequency is slightly lower for sodium iodide than for the other sodium salts in agree- ment with the earlier report on pyridine solutions.10 Solvation band intensities have been observed to be linear functions of concentration in tetrahydrofurang and prOpylenecarbonate.9 Figure 1 shows that the 390 cm'1 solvation band in lithium perchlorate--u-methylpyridine also obeys Beer's law. The addition of tetrabutylammonium bromide to a 0.4M solution of lithium perchlorate in h-methylpyridine results in a gradual shift of the 391 cm"1 band to lower frequency. At a l to 1 mole ratio of lithium ion to the bromide ion the band frequency is 383 cm‘l, identical with that observed Iku'the lithium bromide solution. Further addition of the ‘bromide ion does not affect the frequency of the solvation “band. The results are shown in Figure 2 (the weak band at Figure 1 . 18 Band intensities of lithium perchlorate solutions in h-methylpyridine. D 624 cm'1 perchlorate band. 5535 cm"1 shifted h-methlypyridine band 0 5114 cm.1 Le-methylpyridine band 1 O 390 cm' solvation band. :50 1‘ N ' Ar. .E. A "1’ (\o< a) O V 0 19 360 300 P 240 " chlorate :l K ‘ ‘éuao - dine band E III 1 a: g. 5 “IZO - .- / ./ / ./D --,/.__,o____ 4 / '- -------- 60" .20 o l I l I 0 0.2 04 0.6 0.8 cowc. Limo4 (24) Figure 2. 20 Spectra of lithium perchlorate - tetrabutyl- ammonium bromide solution in u-methylpyridine. 0.1 mm cell 0.433 molar lithium perchlorate ----- 0.417 molar lithium perchlorate + 0.395 molar tetrabulylammonium bromide. 450 21 I - l I ‘\‘ ,\ P‘H \H r’ \ . I | I I \q\ I \ I rabutyl- \‘ I 1w”? T '. I . I l- L‘ I’ a a! I I ‘ I q I r + | ’ \ I e. \ A I, l : : I I , k1 \‘ I \ I, ‘\ I I I t‘ ’1 .‘\ ’ x r" I Q.” I L 450 450 350 3 OO 22 3#3 cm"1 is due to the solvent). The above results indicate that the bromide ion can penetrate the solvation shell of the cation and probably replace a solvent molecule in the inner solvation sphere of the lithium ion. Solvent Vibrational_B§nds Spectra from 600 to 150 cm"1 of #-methylpyridine and its lithium perchlorate solutions are shown in Figure 3. In addition to the solvent bands at 512, #90, and 210 cm"1 the solution spectrum has new bands at 535 and 390 cm'l. The latter is the lithium ion solvation band. The behavior of the 535 cm"1 band is similar to the solvation band in that it varies linearily with the lithium salt concentration (Figure 1). Furthermore, the band undergoes a small shift to 532 and 5#0 cm"1 for lithium bromide and lithium-6 per- chlorate solutions respectively. Figure 1 also includes the salt concentration band intensity plot of the nearby 512 cm"1 solvent vibration. In contrast to the 535 and 390 cm"1 bands, the intensity of the 512 cm"1 band is seen to decrease with increasing salt concentration. 0n the basis of the above evidence it seems reasonable to assume that the 535 cm"1 band arises from the solvent interaction with lithium ion. Further evidence is given by Frank and Rogers who reported that upon complexation by capper chloride, the 512 cm-1 band of #-methylpyridine shifts to 549 cm'1.31 Figures 1 and 2 show that the 535 cm"1 "shifted" solvent band is both more intense and slightly broader than the unperturbed 512 cm'1 band. This 512 cm"1 band is not shifted in sodium salt Figure 3. 23 (a) Spectrum of #-methylpyridine from 575- 200 cm'"1 . (b) Spectrum of 0.31 molar lithium per- chlorate in #-methylpyridine from 600-150 cm"1 cM = 0.1 mm, 3:1mylar beamsplitter, resolution = # cm-l. 21+ L5 f] I I from 573- m per- 600-15-9 05 a: litter: < l V B B A fix I O - l I I 500 400 300 200 cu" .fynulES we: ...-'vJ ‘ I “e ”1' litre. '0‘ 9V 1.: -‘ +““ v. .r. -.:l:du-( aiiition of 251193133 51. d w rlridih‘ :i’m'n q _ "v V‘ 1 .. ‘ 8 § ‘5‘! VII " .“'V%O:~' “Cit-er :E~:f‘“ We. .‘l‘n. - l‘ \ “ME: 0 v L", h‘\ A ‘\ ¢ 7‘. "*e h: Q ‘\ 6 .L a ‘. .73: 25 #-methylpyridine solutions. However, it does have some asymmetry on the high frequency side which indicates that the #-methylpyridine interaction with sodium ion is much weaker than that with lithium. Changes in the vibrational spectra of the solvent molecules were observed in the near infrared and Raman spec- tra for lithium salt solutions in pyridine and #-methylpyridine or in mixtures of the pyridine and nitromethane. Upon addition of lithium salts new bands appeared on the high frequency side of several of the pyridine fundamental vibra- tions. The intensities of these new bands increased with the lithium salt concentration with a concomitant decrease in the intensity of the original band. This concentration dependence of the bands indicates that upon complexation with lithium ion, the fundamental solvent vibrations shift to higher frequency. The limiting shifts for several of the pyridine and #-methylpyridine fundamental vibrations are given in Table 3. No change was observed in the intense 919 cm'1 nitromethane Raman band frequency. Comparison of the spectral data for pyridine and #- methylpyridine solutions indicates that similar vibrational changes occur in both solvents upon addition of lithium salts. All of the Raman and infrared active bands listed in Table 3 are of a1 symmetry and are fundamental planar ring vibrations.29 The Raman data confirm the infrared data reported earlier10 and further demonstrate a shift in the pyridine v vibrations. 12 10 McKinney correlated these vibrational changes of pyridine in lithium salt solutions with those observed by Takahashi, - . » j;e3. solvent lithium pyridin a. "l ..ra'..on U" .K‘ ‘--..a:re’1 fj +‘_1 ‘:‘A . H 0&1.‘ Table 3. Solvent vibrational frequency (cm'l) changes for lithium salt solutions in #-methylpyridine and pyridine-nitromethane solutions. Vibrationa v 8a v 1 v 12 v 6a Infrared #-MePy 1605 993 800 51# +Li+ 1617 1011 806 537 Raman #-MePy 997 803 517 +1.1+ 1014 812 533 Infrared Py 1581 991 603 +Li+ 1597 1003 620 Raman Py 993 1032 +Ll+ 1008 1039 aBand assignments taken from reference 29. fl szlvents. These zeital, planar r1 51:55: frequenci‘ tried to protiC :0 51328.31 ac“ salvation band i: eel: concentratiu The perchlc: relatively trance. 7'3 perchlorate slutions in by:- 1::‘71‘1: perchl or :33 ran “*0 or ‘4‘»! t 331+ ..-l “ band at abs .. . :‘c y. :‘.35 “V Jon“ ‘nitrome: J “!h pair .J '1 c 4. a; Chara ‘ 5am “A 27 33 21.30 for pyridine interacting with hydrogen bonding solvents. These latter authors observed that the funda- mental, planar ring vibrations of pyridine are shifted to higher frequencies when the pyridine molecules were hydrogen bonded to protic solvents. No Raman activity was observed for the lithium ion solvation band in either pyridine or #-methylpyridine at salt concentrations up to 1.5 molar. The perchlorate anion is generally considered to be a relatively noncoordinating ion. The 933 cm'1 Raman band of the perchlorate ion remains unchanged in lithium perchlorate solutions in pyridine or #-methylpyridine. However, in lithium perchlorate pyridine-nitromethane mixtures, when the mole ratio of pyridine to lithium becomes approximately equal to or less than #, a shoulder appears on the 933 cm"1 band at about 939 em'l. Similar behavior of the perchlorate Raman band has been observed in lithium perchlorate- acetone-nitromethane mixtures. This change in the perchlorate ion vibration indicates that the symmetry of the tetrahedral perchlorate ion has been lowered. This decreased perchlorate symmetry when the pyridine to lithium ion ratio is< 1+ indicates that the solvation number of lithium ion is # and, that if the system does not contain enough primary solvating molecules to maintain the solvation shell, the deficiency is made up by the perchlorate ion, which then forms a con- tact ion pair with the cation. In contrast to the appearance of a shoulder on the 512 cm"1 band of #-methy1pyridine as observed in the infrared . R A ww- ‘jf 1.14.» t - y. I. '1’- .4°‘f. 6 Y: :.v..AAA V “3% 33% Q t * ::e:‘.ra in c. ': 4‘“ e. .ect o .v 1H?“ Q‘Qc‘u U vv b- ‘n‘l‘crie 0...».‘0 barl =~‘- nil“: “~03- . l C). d‘Q A. V Q”! J N, " -.=h .“e _ ls .. V -« as .. V. 5 ‘ ‘~ g.“ / ‘\ h . a “.‘V‘x ‘, ’ ‘JQH ' he. Vu‘ 28 for sodium perchlorate solutions, the Raman spectrum shows a definite peak at 52# cm‘l. This shift of the 512 cm'1 band is smaller for sodium than for lithium ions and indicates that the sodium ion --#-methylpyridine interaction is weaker, but of the same kind as the lithium ion--#-methylpyridine interaction. Other Substituted Pyridines Attempts were made to study alkali metal solvation spectra in other substituted pyridines in order to determine the effect of substituent groups on the alkali metal solvation bands. Unfortunately, salt solubilities decreased drastically upon substitution, especially if pyridine was substituted in the 2 position. For example, solubilities of common alkali and ammonium salts in 2,6-dimethylpyridine were too low to allow spectral measurements. The presence of interfering solvent bands in the #00 cm‘1 region rendered impossible a study of 2-methylpyridine solution spectra. The observed solva- tion bands in some substituted pyridines are listed in Table #. The spectrum of 3-methylpyridine shows a strong \h6a out of plane ring vibrational band at #02 cm-l.29 Upon dissolution of lithium perchlorate, two new hands appear, at #12 cm’1 and a broad band at 39# cm"1 (as a shoulder on #02 cm'1 band). These bands were studied in a 1.96 molar solution of 3-methylpyridine in nitromethane which contained varying amounts of lithium perchlorate. Figure # shows that as the 3-methylpyridine to lithium ion mole ration is reduced, there is a gradual transition from the original #02 band 1 to the #12 and 38# cm' bands. Frank and Rogers31 have pointed out that in the complex Cu(3-Mer)?(C1)2 the #02 . . J 4 l .3,-: ”I ‘.,J. ‘1‘ 4.1 f 0' 5 : "39v¢'«':-'7«'r.. ‘__—-‘u—— -~ '.'~IA-’ - 1» ' . ‘ u 0‘ 1‘ ‘A-.vv“ V A . . iuflflw ~ w ‘ fly. . 29 Table #. Lithium solvation band positions for other substit- uted pyridines. 3-methylpyridine LlClUu 31#-dimethylpyridine LlClOu LiBr 2,#-dimethylpyridine LiClOu 2-chloropyridine LiClO LiBr LiI 6LiClo 6LiI 383 379 360 355 3#0 3uo 373 i H- H- 5 cm I . ... . "“4““ .m;4___...—_-.. __ '— in W... at. f Figure #. Lithium perchlorate spectra in a 3-methyl- 30 pyridine-nitromethane mixture. A. B. C. D. Solvent mixture 1.965 M l3-MePy in MeNOZ Mole Mole Mole Mole ratio 3-MePy/LiClOu ratio 3-MePy/LiClOu - ratio 3-MePy/LiC104 - ratio 3-MePy/LiClOu - = 12.5 (OFFSET) %T—-—) i 450 31 I A a 3-methyl- A 5 -MePy 11153: t 12.5 2 9.77 .12 T 3.58 " at J 450 400 3 50 cu" 32 cm-1 solvent band is shifted to #13 cm'l. It seems reason- able, therefore, to postulate that the 38# cm'1 band is the lithium solvation band and the 412 cm'1 band is the #02 cm'1 band shifted due to the lithium ion-solvent interaction. Further support for these assignments come from the obser- vation that upon lithium-6 isotopic substitution, the 38# cm-1 band disappears and a new band appears as a shoulder on the 1 band. high frequency side of the #12 cm- In 3,#-dimethylpyridine-lithium perchlorate solutions, the solvation band at 383 cm-1 is essentially identical with the solvation band in 3-methylpyridine. The band shifts to a position above #00 cm"1 upon lithium-6 substitution, but its exact position is obscured by the strong solvent band at #2# cm-l. The solvation band of lithium bromide in the same solvent is at 379 i 3 cm'l. In two solvents, 2,#-dimethylpyridine and 2-chlorpyridine, the frequencies of the salvation bands for lithium perchlorate solutions are at 360 and 355 cm'1 respectively. In the case of lithium bromide solutions in the latter solvent, the band shifts to 3#0 cm"1 which presumably indicates the formation of the contact ion pair. The lower frequencies for the lithium solvation band in the above solvents cannot be explained solely by the increase in the mass of the sol- vent molecules. It is evident that the substitution in the 2-position decreases the solvating ability of pyridine. The dissolving ability of the two solvents is, however, very low and other common lithium salts as well as common salts of other alkali cations are found to be essentially insoluble. 33 Pyridine The studies of the infrared spectra of sodium and ammon- ium salts in pyridine were repeated and the results reported 10 were confirmed. However, the lithium salvation band earlier reported to be at #20 cm"1 needs some reevaluation in the light of the above results for #-methylpyridine and the other substituted pyridines. 1 First of all the #20 cm" band in the lithium salt-- pyridine solution shows no anion frequency dependence in that the band position for lithium chloride was only # cm"1 lower than that of lithium perchlorate. Pyridine, with a dielectric constant of 12 has an intermediate polarity with respect to acetone and tetrahydrofuran with respective dielectric constants of 26 and 8. Since anion dependence of the salvation band for halide salts has been observed in both acetone7 and tetrahydrofuran“, as well as in #— methyipyridine one would eXpect similar behavior in pyridine. Isotopic substitution in pyridine--lithium salt solu- tions results in shifts of the #20 cm"1 band in the correct direction, but of the wrong order of magnitude. These data and those for #-methylpyridine are shown in Table 5. Table 5. Pyridine--1ithium ion band mass substitution shifts (cm' ). salt pyridine d5-pyridine #-methy1pyridine 7LiClOu 420 389 391 6LiClo4 425 415 416 3# Simple Hooke's law calculations, assuming a "diatomic" model, indicate that substitution of lithium-6 for lithium-7 should increase the salvation band frequency by about 30 cm-l. However, only a 5 cm"1 shift is observed in the #20 cm"1 band. Substitution of pyridine with d5-pyridine gives a 30'cm'1 negative shift when only a 1 cm’1 shift is eXpected. A further increase in solvent mass with #-methylpyridine results in no further decrease in the salvation band frequency. In dS-pyridine and #-methylpyridine the band shifts an substitution of lithium-7 by lithium-6 (3525 cm’l) are more satisfactory with respect to the calcuLated band shifts (g 30 cm'l). The solvating ability of #-methylpyridine should be only slightly greater than that of pyridine due to the inductive effect of the methyl group in the para position. It might be eXpected that the increased donor strength should also increase the salvation band frequency. This factor, however, may be offset by the increase in the solvent mass. The mass increase due to methyl group substitution, however, cannot account for the shift of the lithium salvation band from 420 to 390 om“1. Measurement of the half band widths of the salvation bands indicates that the band width of the 420 cm"1 band in pyridine-~1ithium salt solutions is only about 15 cm'l, whereas in #-methy1pyridine and tetrahydrofuran, the band width is 50 to 60 cm'l. Finally, the behavior of the far infrared pyridine fundamental vibrations have been studied by Frank and Rogers,31 35 who found that two strong bands at 605 and #05 cm'1 increase in frequency upon complexation to first row transition metals. In contrast to the 605 cm’1 band and the other fundamental 1 band is planar ring vibrations discussed above, the #05 cm- assigned to an out-of-plane ring vibration29 which, in very strong complexes, may shift to as high as ### cm’l. Thus it is not unlikely that the 405 cm-1 pyridine band shifts to 420 cm.1 upon interaction with the lithium ion. Examination of the spectra of several lithium salts in pyridine, shown in Figure 5, reveals a broad band on the low frequency side of the intense #05 cm"1 pyridine band. In lithium perchlorate solutions, this band appears as a shoulder with an estimated maximum at about 393 cm'l. In lithium bromide solutions, the band is shifted down to about 383 cm'1 and is partially resolved. This anion dependence can be demonstrated by the addition of equimalar tetrabutylammonium bromide to a lithium perchlorate-- pyridine solution. The band is shifted to virtually the same position as observed in lithium bromide solutions in pyridine. The lithium-6 perchlorate spectrum in pyridine, (Figure 5) in the #00 cm"1 region appears to be composed of the pyridine #05 cm'1 band and a new band at about #28 cm‘l. This latter band might be rationalized as resulting from the overlap of the #20 cm'1 band with the salvation band shifted to higher frequency. A general purpose curve fitting programBu was used to fit the lithium bromide and lithium-6 perchlorate pyridine spectra. The salvation band in #- methylpyridine was best fit to a gaussian line shape. Figure 5. 36 Spectra of some lithium salts in pyridine, 0.682 M LiClOu __ 0.#96 M LiBr ....... 0.508 m 6LiC10 4 . . . . pyridine ik-——" ABSORBANC E. 37 5‘00 25 moz....<._wm (M) CONC. ## lithium ion in pyridine--nitromethane solutions. The some- what lower value obtained here is not surprising since the dielectric constant of the "inert" solvent, benzene, is much lower than that of nitromethane. The lower dielectric con- stant of the benzene-pyridine mixture should allow greater ion pairing of the lithium and perchlorate ions with the anion replacing one of the solvating pyridine molecules. Conclusion From the above discussion, it is seen that the combin- ation of infrared and Raman spectroscopic techniques are very useful in the study of species present in ionic solutions. It has been shown that pyridine and substituted pyridines solvated lithium, ammonium and sodium salts with the formation of free solvated ions or solvent separated ion paris for salts of polyatomic anions. In the case of halide salts, there is evidence for the formation of contact ion paris. The position of the lithium salvation band in 1 pyridine has been reassigned to about 390 cm' with lower values observed for the halide salts. Thus all of the lithium salvation bands in pyridine and substituted pyri- 1 dines appear at or below 390 cm" and are strongly influ- enced by substitution at the 2-position. NMR HISTORICAL INTRODUCTION AND THEORETICAL Nuclear magnetic resonance (NMR) has become a most useful and powerful technique for the study of chemical systems. The application of high resolution NMR techniques, particularily proton, fluorine-19, and carbon-13 chemical shifts, to organic molecular structure determination is familiar to most chemists. These techniques have been adequately covered in a number of texts.35-39 The chemical shift arises from the fact that a magnetic nucleus may eXperience a variety of local magnetic fields resulting from the surrounding electronic motion as modified by chemi- cal bonding and molecular association. These local fields "seen" by the nuclear magnetic moment are usually small-- 10'"7 to 10'“ of the applied field, but are consequences of the nuclear electronic and, therefore, chemical environment. Atomic chemical shifts may be considered to be a sum of diamagnetic and paramagnetic contributions. The diamagnetic term (Gd) arises from the inner, spherically summetrical electrons which set up a small magnetic field apposed to a, large, externally applied mag— netic field. Dickensonl’O has calculated the order of magni- tude of this contribution to be, 60.: Cu (O) = -3.19 x 10"5 Zu/3 (1) 45 #6 where C is a constant, v(C) is the electrical potential at the nucleus from the core electrons and Z is the atomic number. Jameson and Gutowsky41 have more recently discussed the apparent increase in observed chemical shift with increas- ing atomic number and pointed out that although the diamagne- tic contribution to the chemical shift can be calculated, the often predominating paramagnetic contribution is difficult if not impossible to calculate. This contribution arises from the anisotropic distribution of electrons in the non spherically symmetrical outer electron orbitals. Further- more, for nuclei of atomic number greater than one, the par- tially papulated excited electronic states contribute to the paramagnetic term and add to the complexity of calculating this contribution to the chemical shift. The paramagnetic term may be formulated41 as op = §_q(l.’> (2) AE'r'B where AB is the electronic excitation energy and the l/r3 term is the expectation value for the outer electrons calculated over all of the populated electronic states. Chemical shift measurements have made up the bulk of NMR studies reported in the literature due to the technico- logical advances in and the availability of high resolution NMR spectrometers. However, equally important information, which is complementary to chemical shift data, can be obtained through the study of NMR relaxation times. Because relaxa- tion time measurements have been extensively used in the study of electrolyte solutions and are less well known to most chemists, relaxation processes and information derivable #7 from relaxation time measurements will be reviewed. If we place a magnetic moment (u) in a larger magnetic field (Ho) it will tend to process about the applied field at a rate which is the Larmor frequency and is proportional to H If we assume a collection of nuclei of spin 1/2 O. we find two quantum states populated according to the Boltz- man distribution equation and the net magnetization of each state is aligned either parallel or antiparallel to the applied field, H The states are popuLated according o' to the familiar Boltzman distribution NZ/Nl = exp (-AE/kT) where the separation of the energy levels is AB = 2nH.35b Thus in a strong magnetic field there is a slight excess of nuclei in the lower energy state with a net resultant mag- netization along Ho. Transitions between the lower and upper states can be induced by the application of radio frequency (rf) radiation at the Larmor frequency with its concommitant magnetic component (H1) perpendicular to Ho. We may further simplify the system by performing a rotation of the coordinate axes to coincide with the pre- cessing magnetic moment. Thus it would appear to the nucleus that the applied Ho is processing about it at the “armor frequency. “y this conversion to the "rotating frame"42a we "freeze out" the precession and consider only changes of the resultant magnetization. Now if we induce transitions from the lower to upper magnetic states, we will eventually equally populate both states. No more energy can be absorbed nor will there be any net nuclear magnetization. The fact that we observe the absorption of energy implies that there #8 is a mechanism fortinon-radiative relaxation of the absorbed energy to return the system to magnetic equilibrium. This process, known as spin lattice or longitudinal relaxation, occurs with a time constant, T1, by transferring the excess energy to nearby solvent molecules or ions collectively known as the lattice. The mechanism arises from electromagnetic field gradients associated with thermal motions of lattice components which may produce transitory fields on the order of the Larmor frequency and permit energy transfer to occur. The spin lattice relaxation time determines primarily the rate at which a disturbed spin system returns to magnetic equilibrium and is observed in the laboratory as being re- sponsible for resonance line saturation (equalization of the spin states populations). The spin-spin relaxation time, T2, is related to the phase coherence or "memory" of a displaced resultant mag- netization. This may be envisioned by tipping a collection of magnetic nuclei in the rotating frame 90° into the x axis from their equilibrium value in the z axis. Assuming T1 is long compared to the eXperiment time we may observe a de- crease in the resultant x axis magnetization as the individual nuclear moments "fan out" in the xy plane. The decrease in the x axis resultant magnetization or loss of phase coherence is described by the time constant T2. The mechanism for T2 relaxation is the interaction with other magnetic nuclei through local and external magnetic field inhomogenieties. Spin-spin relaxation does not involve #9 net changes in energy level population, but is indicative of the time a particular nucleus remains in the excited state. The "natural" linewidth of a nuclear resonance is inversely proportional to T2. The major contributions to T2 in solutions are the magnetic field inhomogenities and viscosity effects. T2 measurements from linewidth deter- minations are valid only in the case where the resolution HO 1/2, the observed linewidth at half height.“23 ‘ 1! v = l = _1__ + (9113/2) (3) 1/2 ‘T2; T2 0 Hahn43 in his classic paper described the spin echo technique for T2 measurement. After the 90° shift of the equilibrium resultant moment at the time T he applied an rf pulse which effectively rotated the coordinate system 180° about the y axis. At time 21, the magnetic moments refocus along the -x axis producing an echo. A plot of echo amplitude versus time then yields T2. Important modif- ications of Hahn's T2 measurements have been made by Carr and Purcell,“l and also by Meiboom and Gill.u5 A review of T1 and T2 relaxation time measurements can be found in chapter 2 of reference #2. The existence of spin-lattice and spin-spin relaxation fimes would be merely ancillary to chemical shift studies with- out correlation to molecular and ionic processes. We may de- fine Tc' as the "average time between molecular collisions “2b The inverse for a molecule in some states of motion." correlation time (l/Tc) may be described as the upper limit of fluctional electromagnetic frequencies, which are observed at a nucleus and are derived from thermal motions of inter and intra 50 molecular or ionic lattice components. It follows that To should be a function of temperature and viscosity. Bloem- bergen, 23 gl.“6 applied Debye's theory of dielectric relaxa- tion and found that the correlation time was a function of viscosity and temperature, chi n/T. They noted that at small values of To, for nonviscous fluids, T1 and T2 are both proportional to To' At very low values of Tc only Tl became inversely proportional to T0 with the minimum in the T1 vs. T curve responding to the internuclear fluc- c tional frequency, l/Tc =‘”ov the Larmor frequency. There- fore T1 processes are related to the fluctional electro- magnetic frequencies at the Larmor frequency, while T2 pro- cesses are functions of the correlation time itself. Now then, we can relate processes which produce fre- quencies at the Larmor frequency to T1. A list of such processes which provide coupling between the observed nucleus and the lattice is given in Table 6. I Table 6. T1 related processes.L2b 1. Magnetic dipole- dipole interaction 2. Electric quadrupole interaction 3. Chemical shift anisotropy interaction #. Scalar - coupling interaction 5. Spin - rotation interaction The general functional relationship between these processes is of the form -1 _ 2 , T1 — EC f(rc) (4) rePT95 .‘rere i- i. c frictional d8?“ 1:: are reviewF * contrib- A \) h ‘3. :he strength at action 3:: be elation #. 51 , #2b where he represents the strength of interaction. The functional dependencies are given by Farrar and Becker”2 and are reviewed in detail by Hertz.47 Therefore, with T1 and Tc contributions known from a specific interaction, the strength and specific molecular details of the inter- action may be deduced from the specific dependence of equation #. EVOLUTION OF NMR TECHNIQUES Before the nuclear resonance effect had been observed in solids and liquids, the nature and existence of discrete magnetic states in many nuclei were known from emission spectrographic and molecular beam studies. In 1936 and l9#2 #8,#9 Gortner reported attempts to observe induced transitions between the magnetic states of lithium-7 nuclei in lithium fluoride. both attempts failed due to use of "unfavorable" samples.36a BlochSO and Purcell51 in l9#6 simultaneously reported observing hydrogen resonance absorption in water and paraffin wax respectively. Their success was rewarded with a joint Nobel prize in 1952. With the theoretical basis of BlochSZ, early NMR studies were concentrated on the study of T1 (spin-lattice) and less rigorously on the study of T2, (spin-spin) relaxation times. Rollin, 22 £1.53 observed the lithium-7 resonance of lithium fluoride down to 2°K. Bloembergen, Purcell, and Pound,“6 in l9#8 measured T for water and other protonic liquids and correlated the 1 dependence of T1 and T with the solution viscosity for 2 glycerin protons. They also studied and theoretically de- scribed the effect of paramagnetic ions on the water proton Tl relaxation and further compared the observed T changes of l protons and lithium-7 ions in paramagnetic solutions. 52 53 Knight54 in l9#9 first reported the phenomenon subse- quently to be known as the chemical shift. Dickenson55 in 1951 reported a NMR survey of several nuclei in liquids or aqueous salt solutions. He observed chemical shifts for boron-ll, nitrogen-l#, flourine-l9, and phosphorus-31. Because the smallest chemical shift he could detect was 5 ppm, he observed resonances for but did not observe chemical shifts for the proton, lithium-7, sodium-23, and aluminum- 27 nuclei. With the above-mentioned ground work of Bloembergen, 23 51.46 and Dickenson55 the NMR study of electrolyte solu- tions grew dramatically. The great majority of such studies have been directed toward the elucidation of liquid struc- ture, ion-solvent and ion-ion interactions in water. The great amount of literature has fortunately been reviewed 56-60 extensively by several authors. WATER PROTON CHEMICAL SHIFTS The structure of pure water has been studied in detail by Hindman61 who observed water proton chemical shifts as a function of temperature. He found a nearly linear 1 ppm shift over the temperature range 0 to lOOfi C which was cor- related to the breaking of H-bonds with increasing tempera- ture. Water proton shifts in electrolyte solutions have been extensively studied. Shoolery and Adler62 interpreted observed proton chemical shifts as resulting from individual ionic contributions to increased or decreased water structure. 5# These studies were extended by Hindman63 who considered var- ious contributions to the chemical shift. He concluded that of the alkali metal ions only lithium ion gave any indication of ordering (structure making) of water molecules. Glick, _£_§l.6u also observed aqueous proton shifts in alkali halide solutions and related the upfield shits to increased hydrogen bonding and downfield shifts to increasing anion size and a decrease in hydrogen bonding. They further correlated the NMR results with the infrared and Raman data. One of the major difficulties with the determination of the relatively small chemical shifts (1 ppm) observed for water protons in aqueous electrolyte solutions is that of proper referencing. Suitable internal standards are preferred, but must not alter the chemical system by their presence. Gordon and Thorne65 examined 22 different internal references and found varying degrees of interaction between the internal reference material and the water protons. Davies, 23 91.66 have reviewed the effect of electrolytes on water proton chemical shifts and attempted to set up a chemical shift scale based on the assignment of an absolute chloride ion effect on the proton chemical shift. These authors concluded that although long rnage water structural effects may be present, consideration of inner sphere co- ordination of water by various ions can adequately explain the aqueous proton shifts. The temperature dependence of the water proton chemical shift in aqueous alkali salt solutions was used by Nalinowski, gt al.67 to determine the total salt hydration numbers. They sreeulated tha‘ 2&1?th effects 1: tetraalkyla: 3201926 C anc :uemkers an: seeing as the trrcugh the te‘ In connec‘ theiindings o: :55 :agnetic 31 21:91.. A line “:3"- increase( 5 and the 1...“: was r915 55 speculated that at salt concentrations below 5 molar, ion pairing effects were observed. Water proton chemical shifts in tetraalkylammonium halide solutions have been studied.68 Above 26 C and at about 0 C the cation acted as water struc- ture makers and breakers respectively with the effects in- creasing as the cation size increased from tetramethyl- through the tetrabutylammonium ion. In connection with these NMR water structure studies, the findings of Ergin and Kostrova69 should be mentioned. The magnetic susceptibility temperature dependence of alkali metal halide salts and their aqueous solutions were deter- - x was found mined. A linear difference, Ax= X cryst soln which increased with atmoic number for both the alkali metals and the halides. The temperature at which Xsoln = Xcryst was related to the individual ionic contribution to the aqueous solution structure. NMR Relaxation Studies In Aqueous Electrolyte Solutions Beginning with the work of Bloembergen, Purcell, and 46 Pound mentioned earlier, a great amount of work has been done using NMR relaxation techniques to study aqueous elec- trolyte solutions. This work has been reviewed by Deverell58 and Hertz.59'7o The effect of diamagnetic salts on the water proton T1 in aqueous solutions correlates well with the chemical shift results mentioned above. Highly electropositive cations such as magnesium, calcium and lithium increase the water proton relaxation rate (l/Tl) relative to pure water because of in- creased interaction thraugh water structuring. Likewise, 56 large, singly charged cations and anions disrupt the water structure, which results in lower water proton relaxation rates.58 The two most important relaxation mechanisms encountered in electrolyte solutions are magnetic dipole-dipole and quad- rupole. The latter generally predominates in cases where a nucleus with a quadrupole moment is involved. Fabricand and Goldberg71 studied the water proton relaxation rates in lithium-7 and lithium-6 chloride solutions. Since the lithium-6 isotope has a much smaller magnetic moment than lithium-7, the difference between the water proton relaxa- tion rates for the two isotopic solutions is due to the greater proton-lithium-7 dipole-dipole interaction. From the net dipole-dipole relaxation rate, they concluded that the correlation time for water molecules in the lithium ion inner salvation sphere is no longer than about 10"11 seconds. Some of the strengths and weaknesses of the use of NMR relaxation studies to determine water structure have been discussed.72 However, Arnold and Packer73 have presented some questions about the validity of determination of correla- tion times through the viscosity of systems in which electro- lytes strongly effect the solvent structure, i.e. water. Larsen?“ has reported correlation times and reorientation activation energies for symmetrical and unsymmetrical tetra- alkyammonium cations in aqueous solution. He observed the alkyl proton relaxation times which he found to be dependent on the rate of quadrupole relaxation of the nitrogen-l4 57 nucleus as determined by the ion tumbling rate. The reorien- tation activation energy was found to be low for symmetrical ions and to increase as one of the alkyl groups was length- ened. He further studied the nitrogen-l4 quadrupole governed proton lineshapes and T1 relaxation rates of pyridine protons in lithium halide salts dissolved in pyridine-water and neat pyridine solutions.75 From an anisotropic rotational diffusion mechanism he proposed a linear lithiumixnbu- chloride-"pyridine species in lithium chloride solutions in 5 percent pyridine-water. For lithium bromide solutions in the same solvent mixture, he found that a water molecule replaced the halide ion in the linear complex. In neat pyridine-lithium chloride solutions his data were rationalized by a tetrahedral arrangement of 3 pyridine molecules and a chloride ion about the lithium ion. Takuhiro, 23 £1.76 have studied the structure of water dimethylsulfoxide mixtures by determination of T1 relaxation times of the water and dimethylsulfoxide protons. They found maxima in the relaxation rate versus composition curves at about 0.65 mole fraction water, which indicates that the solution undergoes various structure changes with composition. Variation of temperature, and the use of deuterated solvents to dilute the intermolecular magnetic effects gave further information about the nature of the structural species. In contrast, the proton chemical shift results showed only smooth transitions between the two extreme compositions of pure water and dimethylsulfoxide. 58 RELAXATION STUDIES WITH "OTHER NUCLEI" It has been mentioned above that quadrupolar nuclear relaxation processes usually predominate in the cases where the relaxing nucleus has a quadrupole moment. However, the lithium-7 ion with a nuclear spin of 3/2 and a significant quadrupole moment of -0.0# x 10-24 cm2,77 has been found by Woessner. £3 £1.78 to undergo nuclear relaxation by both quadrupole and magnetic dipole-dipole mechanisms in aqueous lithium salt solutions. The unexpected by low quadrupolar contri- bution to the relaxation rate of lithium-7 ion indicates that the ion is in a highly symmetrical environment. The above authors separated the dipole-dipole relaxation rates from the quadrupolar rates by determining the lithium-7 relaxa- tion rate (l/Tl) for lithium chloride in water and deuterium oxide. The lithium-7 relaxation rate in water was found to be about twice that observed in the deuterium oxide solutions. Because deuterium has a much smaller nuclear magnetic moment than the proton, the difference in relaxation rates in the two solvents was attributed to the greater nuclear dipole- dipole interaction between lithium-7 and water protons. Although the two relaxation mechanisms contribute equally to the lithium-7 relaxation in aqueous solutions, the shape of the relaxation rate versus salt concentration curves was found to be determined by the quadrupolar relaxation process. The curves for lithium chloride solutions show a break at concentration of 4.4 molar which indicates that the lithium ion undergoes a change in environmental symmetry at this concentration. 59 Hertz, 22 91-79 studied the relaxation rates of lithium-7 chloride, bromide and iodide in water and deuterium oxide. They determined from the lithium-7--water dipole-dipole interaction that the water molecules remain in the inner salvation sphere long enough to observe rotation effects of the water molecule about the lithium-oxygen axis. The quadrupolar contribution to the relaxation rate was related to the break down of cubic symmetry in the lithium ion inner hydration sphere as the salt concentration increases. Mishustin and Sidorova8O determined the lithium—7 spin-lattice relaxation time for several lithium- salts in aqueous solutions and observed a strong concentration depen- dencefxu'lithium salts of weak acids. The concentration dependence was interpreted in tenns of solution viscosity changes. 81 studied lithium-7 relaxation times Craig and Richards of lithium chloride solutions in water, methanol, formic acid and dimethylformamide. The first three solvents gave almost identical relaxation rates when corrected for visco- sity differences. The relaxation rate for dimethylformamide, however, was about five times greater than that for other solvents which indicates a much stronger lithium ion-solvent interaction. The lithium-7 chemical shifts in the above solvents were also determined, but were found to be too small to be measured accurately. Lithium-7 spin-lattice relaxation times and chlorine-35 linewidths were determined by Bryant82 in concentrated aqueous lithium chloride solutions. The viscosity adjusted 60 relaxation rate for both ionic nuclei showed marked concen- tration effects when the water to lithium chloride ratio was about 12, 6, 5, or 4 to 1. Chizhik and Ermakov83 studied lithium-7 and sodium-23 aqueous chloride relaxation times as functions of concentration and temperature. The sodium- 23 relaxation time decreased with increasing temperature, whereas that for lithium-7 gave a distinct inflection in the 25-30° region. As reported earlier77 the relaxation rate of lithium-7 was found to decrease as the solvent was changed from water to deuterium oxide. However, the authors re- ported that the sodium-23 relaxation rate increases with the same solvent change. Hertz, 23 31.70 have reported relaxation studies of lithium-7, sodium-23, and rubidium-85, bromide-81 and iodide-127 salts in aqueous solutions. They concluded that there is no direct cation-anion contact for lithium iodide and bromide as well as sodium iodide at salt concen- trations below seven molar. However, such ion-ion contact is observed for the rubidium salts. The authors were un- able to correlate the relaxation data with the iodide-127 chemical shifts in various alkali and alkaline earth iodide salt solutions. Halogen NMR studies of electrolyte solutions have recently been reviewed by Hall.84 Alkali Metal Ion Chemical Shifts in Electrolyte Solutions The chemical shift of alkali metal cations in aqueous solutions has recently been reviewed.57'58 The magnitude of the chemical shifts increases with the atomic number as 61 discussed in the introduction to this section. Deverell and Richards85 studied the concentration dependence of the chemical shift for sodium—23, potassium-39, rubidium—87 and cesium-133 in aqueous solution. They found that the alkali metal ion chemical shift was dependent on the salt concentration and varied with the anion. The cation shield- ing by the anion increased in the order iodide, bromide, chloride, fluoride and nitrate. The anion effects were observed to be smaller for the sodium ion than for the other alkali ion studies, which was attributed to the higher degree of salvation of the sodium ion. Akitt and Downs86 reported small, but similar, anion effects for lithium-7 ion aqueous solutions. However, they later recanted their claim due to an error in their diamagnetic succeptibility corrections.87 Lutz88 accurately determined the magnetic moments of lithium-6, lithium-7, sodium-23, rubidium-87 and cesium-133 ions in aqueous solution. The magnetic moments of the latter two nuclei were found to differ from their gas- eous state moments as determined by other methods because of the magnetic shielding of the ions by the solvent. Sodium—23 NMR has been used to study salvation effects of sodium ion in non-aqueous solution. Bloor and Kidd89 studied the influence of solvent as well as the salt concen- tration dependence of the sodium-23 chemical shift in several solvents. Erlich, st 21.90 determined sodium-23 chemical shifts for sodium iodide and tetraphenylborate in various solvents. In most non-aqueous solvents they observed a downfield shift of the sodium-23 resonance in solutions of 62 the iodide with respect to the tetraphenylborate, which indicated that ion pairing effects were being observed. A plot of the chemical shift in sodium tetraphenylborate solu- tions versus Gutmann's solvent "donor numbers"91 gave a linear relationship. The solvent "donor number" is simply the enthalpy (in kcal/mole) of complex formation between the given solvent and antimony pentachloride in 1,2- dichloroethane solutions. The relationship between the sodium chemical shift and Gutman's empirical solvent donor scale was used by Herlem and Popov92 to determine the "donor 93 numbers" of several amine bases. Erlich and Popov ex- tended the work reported above90 to include sodium thiocyanate and perchlorate salts in several solvents. They found that for the tetraphenylborate and perchlorate cases, the sodium- 23 chemical shift was independent of the salt concentration. whereas that for iodide and to a lesser degree thiocyanate, it was concentration dependent. Erlich, gt al.9u have applied sodium-23 NMR to the study of several binary nonaqueous solvent mixtures. Their study of sodium tetraphenylborate solutions in pyridine DMSO mixtures indicated that the di- methyl sulfoxide preferentially solvates sodium ion even though the donor numbers of the two solvents would indicate that pyridine should be the slightly better solvating agent. The NMR results were confirmed by observing the sodium ion- 1 as the sol- vibration frequency change from 200 to 180 cm' vent composition was varied from DMSO to pyridine. The change in frequency parallels that of the sodium-23 chemical shift which again indicates preferential salvation of sodium ion 63 by dimethylsulfoxide. These results were interpreted as resulting from the disruption of the dimethylsulfoxide solvent structure by pyridine. Lithium-7 chemical shifts in lithium salt-nonaqueous solvent systems have not received a great deal of study. To date only two original references have been reported. haciel, g3 £35 and Akitt and Down887 reported lithium-7 chemical shifts for perchlorate and bromide salts in several nonaqueous solvents. They observed that the total chemical shift range is about 6 ppm and within this range the chemical shift is solvent dependent. These solvent dependencies were not correlated with solvent molecular properties such as the dielectric constant, dipole moment or heat of vaporiza- tion, but could be rationalized to some degree by the specific solvent effects. For instance, the carbonyl-containing solvents were grouped together, which indicated that the lithium ion is most influenced by the nature of the solvent interaction site. This specific solvent effect was further pointed out by the extreme lithium-7 ion chemical shifts in pyridine and acetonitrile. These shifts respectively parallel the low and high field shifts observed for aromatic and acety- lenic protons. The chemical shift was also dependent on the anion present in solution as shown by the fact that the lithium-7 chemical shift in lithium bromide solu- tions was at slightly lower field than in the lithium perchlorate solutions. Lithium-7 NMR has been used to study the complexation of the lithium ion with nitrilotriacetic acid96 (HBNTA). 64 The data were intrepreted as resulting from the formation of a Li(NTA)2+5 complex. Lithium-7 and Lithium-6 NMR techniques were used by Attalla and Eckstein to determine the isotopic ratios in isotopic mixtures.97 From the above discussion it is seen that lithium-7 NMR is a sensitive probe into the nature of electroylte solution. The lithium-7 nucleus is well suited to high resolution techniques because of its narrow linewidth (generally less than 0.3Hz) and its good sensitivity (0.29 with respect to proton = 1.0 at constant field). NMR EXPERIMENTAL Solvents Pyridine: reagent grade pyridine, was refluxed for twelve hours over barium oxide and fractionally distilled through a 100 cm glass helicies packed column. The product was stored over barium oxide in an amber bottle. 4-methylpyridine: this solvent was obtained from the Aldrich Chemical Company and was purified and stored in the same manner as pyridine. Other Substituted Pyridines: the other substituted pyridines. obtained from the Aldrich Chemical Company, were refluxed over barium oxide for a minimum of two hours and fractionally distilled from barium oxide. The distillation products were stored over barium oxide. Acetonitrile: acetonitrile, Fisher, certified reagent grade, was refluxed over P205 for two hours and fractionally distilled. The product was then refluxed over barium oxide for two hours and again fractionally distilled and stored over Fisher type #A molecular sieves. Acetone: reagent grade acetone was dried over calcum sulfate for twenty-four hours, then fractionally distilled from fresh calcum sulfate. The product was stored in an amber bottle and kept in a dry nitrogen atmosphere. 65 66 Acetic Acid: acetic acid, Fisher ACS reagent grade, was refluxed overnight over phosphorus pentoxide and fractionally frozen twice, M.P. 16.7° C, and stored in an amber bottle in a dry box. Dimethylsulfoxide: reagent grade, J. T. Baker, dimethylsul- foxide was dried over Linde type #A molecular sieves and vacuum distilled at 50° C at 0.1 mm pressure. Karl Fisher titration gave 0.051 water. Nitramethane: Matheson, Coleman, and Bell practical grade nitromethane was purified by passing it through a 30 cm column of Dowex 50w-X8 cation exchange resin and drying over Drierite for several hours. The solvent was then distilled from fresh Drierite through a 50 chVigreux column. The fraction boil- ing at 100.2° C was taken and stored in an amber bottle. The water concentration was found to be 0.017 molar by Karl Fisher titration. Tetramethylguanadine: tetramethylguanadine, obtained from American Cyanamid, was refluxed over barium oxide for twenty-four hours followed by vacuum distillation at 36 to 38° C and 0.1 mm pressure. Lithium Salts: reagent grade salts were dried at 180° C for three days. The dried salts were stored in a dessicator aver magnesium perchlorate and kept inma dry box. Tetrabutylammonium Salts: tetrabutylammonium perchlorate and bromide were dried for forty-eight hours at 60° C in vacuum and stored in a dry atmosphere. 67 SAMPLE PREPARATION Due to the sensitivity of the lithium chemical shift to the presence of water, the samples were prepared under as nearly anhydrous conditions as possible. Most of the sol- utions were prepared entirely in a dry box with a dry nitro- gen atmosphere. The salts were weighed on a Wetler type H6 analytical balance within the dry box. Problems with the air pressure changes affecting the balance were reduced by using an air ballast in the form of a sealed glove bag partially filled with dry nitrogen connected to the dry box. An inter- val of at least thirty seconds was allowed for the pressure to equilibrate before final weights were recorded. The salts were weighed directly into a volumetric flask and made up to the mark with solvent. The solutions were transfered to standard 5 mm NMR tubes, a reference capillary inserted, and the tube sealed with a NMRS 15-105 NMR tube pressure cap. The reference solution was contained in a sealed capillary tube, 1.6-1.8 x 90 mm (Kimax, number 34505) and held centered in the NMR tube by two 3 x 4 mm Delrin washers, one at each end. INSTRUMENTATION Two instruments were used to obtain the lithium-7 NMR data. A Varian DA-60-IL was used in its standard frequency lock configuration and operated at a Larmor frequency of 23.315 MHz. and a magnetic field of 14.1: Kilogauss. The field was locked to the reference signal and the sample peak recorded. Positioning the recorder pen at the peak 68 maximum and counting the sweep and lock oscillators with a Monsanto Model 100B frequency counter gave by difference, the chemical shift with a precision of: 0.1 Hz. The magnetic field was tuned with the reference sample, usually 4 molar lithium perchlorate in water to the homogeneity limit of the magnet-~0.2-0.3 Hz. In almost all cases the magnetic field inhomogeniety limited the observed linewidths. The spectra were obtained at ambient probe temperatures, about 28° C. Temperature control for the temperature studies was obtained using a Varian Vh343 temperature controller. Lithium-7 chemical shifts were also obtained at 60 MHz on either a NMR Specialties MP-lOOO operated in the "time- sharing" mode or a spectrometer built of commercial com- ponents and assembled in this laboratory. The latter instru- ment used a General Radio 1165-AR7, 0.01=70 MHz frequency synthesizer with the frequency sweep controlled by a Fabri- Tek 1082 computer. The computer was then used to time average the spectra. The synthesizer output was pulsed at 50 kHz with signal detection by the single sideband suppressed carrier method with a phase sensitive detector. The 5 mm crossed coil probes were operated at 60 MHz. The magnetic field was derived from a Westinghouse 6O kilogauss niobium base alloy superconducting solenoid. Further homogenization of the field was provided by normal shim coils concentric to the probe. For lithium, the magnet was operated at about 36.2 kilogauss with the rf frequency of about 60 MHz. A 97 MHz probe was constructed, but insufficient development time precluded its use to take advantage of the full field capability of the magnet for lithium. The best 69 resolution observed using the "time-sharing" or psuedo continuous wave methods with this magent was 3H2. A few spectra were obtained with the MP 1000 in the pulse mode using the 1082 computer, and a fast Fourier transform- program. The results were promising but unfamiliarity with the system degraded the results. NMR RESULTS AND DISCUSSION The study of lithium-7 NMR was begun with a study of possible reference solutions. Aqueous salt solutions have m; ‘ - . V» g A=-—-;.‘2‘~ been generally used as alkali metal NMR references. Lithium-7 chemical shifts in lithium chloride and perchlorate aqueous snail-"w. “m. 1.; In ~ 272': solutions were determined and the results are shown in W" .1 Table'7. The small but significant downfield shift with increasing salt concentration is similar to that found earlier,85'86 which was related to the differences in the bulk diamagnetic susceptibility of the concentrated salt solutions. The reference solution adopted for this study was four molar aqueous lithium perchlorate. Because no lithium compound was available for use as a lithium-7 internal standard, all chemical shifts were measured against a reference salt solution isolated from the sample solution. The difference between the diamagnetic susceptabil- ities of these solutions leads to an apparent chemical shift for which allowance must be made.360 Corrections may be applied from the known36d or calculatedS%3diamagnetic susceptabilities (x\) of the sample and reference through the equation 6 corr = (gobs + 2/3" (kref '16:) (5) which applies to concentric cylindrical sample tubes, per- pendicular to the applied field. However, when a supercon- ducting solenoid is the source of the magnetic field, the 70 71 Table 7h Observed lithium-7 chemical shifts of aqueous saltsa. 0 concentration concentration ca. 15 fl ca. 4.5 fl LiCl LiClOA Conc. <5Hz 622m Conc. 6 Hz 6 ppm satb 33.6 1.uu sat° 46.0 1.97 sat/2 39.4 1.69 4.0 M 45.7 1.96 sat/4 41.2 1.77 2.0 44.6 1.91 sat/8 42.4 1.81 1.0 44.0 1.89 a reference = 0.6 fl LiClOu in pyridine b « “n. *‘l! . . .- I' " '0 '.. u“ r n .y . 72 applied magnetic field is parallel to the sample tube and 99 the diamagnetic correction has been shown to be given by the expression 6 =... r8f_ + corr ”/3 fl 0% xv) 6obs (6) The corrections of the chemical shifts are given in Table8 and are seen to vary from 0 to 1 ppm depending on the magnet used and on the solvent diamagnetic susceptibility. 4" f“ 71". _ t! The corrections were calculated only for the solvent diamag— r‘ \ fi’"—,‘ netic susceptibility. Because the diamagnetic susceptibility of water changes only about 47 with the addition of 4 molar lithium perchlorate and since most of the solutions studied were less than 0.5 molar, the salt contribution to the bulk diamagnetic susceptibility was considered‘to be negligible. Lithium-7 chemical shifts of lithium perchlorate and bromide salts in various solvents obtained at 59.6 MHz are listed in Table €9and are shown in Figure 8. In acetonitrile solution the lithium perchlorate chemical shift varies linearly with salt concentration. However, the chemical shift for the bromide salt at concentrations greater than about 0.15 molar are about 1 ppm downfield of that for lithium perchlorate. At lower concentrations, the chemical shift increases and approaches that for lithium perchlorate. This behavior of the lithium-7 chemical shift in acetonitrile—- lithium bromide solutions indicates that the lithium ion experiences at least two different chemical environments. The magnitude of the chemical shift difference between the two salts above 0.15 molar concentrations suggests that lithium bromide contact ion pairs are formed in lithium _ 73 Table 8. Diamagnetic susceptibility corrections for various solvents. Xv (4120" Xsolv) “cox-r EM 6corr so Solvent -1x106 -1 x 10 6 -lx106 -1x106 acetone 0.460 0.261 0.546 1.093 acetic acid 0.552 0.169 0.354 0.708 acetonitrite 0.534 0.187 0.392 0.783 methanol 0.515 0.206 0.431 0.863 pyridine 0.612 0.109 0.228 0.4565 4-methy1pyridine 0.614 0.107 0.224 0.448 nitromethane 0.346 0.784 0.784 1.568 tetrahydrofuran 0.641* 0.079 0.165 0.331 dimethylsulfoxide 0.005*4 0.115 0.241 0.481 *estimated values 1.... .__..--_,_A _ __._--. ‘1 . _ 31' 'l 74 Table 9. Lithium-7 chemical shifts in various solvents at Acetonitrile Conc ' obs . Gobs . 6 6c orr 6 fl Hz. x10 x10 LiClOu 0.573 +105 1.76 2.545 E 0.428 104 1.745 2.53 i 0.228 103 1.73 2.51 0.0855 99 1.65 2.43 7 0.057 99 1.65 2.43 0.028 100 1.68 2.46 LiBr 0.503 44 0.74 1.52 0.376 50 0.84 1.62 0.251 48 0.805 1.59 0.126 53 0.89 1.67 0.075 56 0.91 1.72 0.037 59 0-99 1.77 0.025 64 1.09 1.87 Acetone LiClOu 0.516 -134 -2.24 -1.15 0.387 -137 -2.30 -1.21 0.258 —137 -2.30 -1.21 0.0775 -142 -2.38 -1.29 0.052 -l46 -2.45 -1.36 Table 9. (con't.) 0.418 -l53 0.314 -155 0.209 -l55 0.105 —157 0.063 -156 0.042 -156 0.021 -155 Acetic Acid 0.468 - 27 0.350 - 30 0.117 - 30 0.071 - 30 0.047 - 27 0.023 - 27 0.396 - 41 0.297 - 40 0.198 - 43 0.099 - 43 0.060 - 41 0.040 - 42 0.020 - 43 Dimethylsulfoxide 0.526 35 0.394 34 —2.56 -2.60 ~2.60 —2.63 —2.62 -2.62 -2.60 +0.26 +0.21 +0.21 +0.21 +0.26 +0.26 +0.02 +0.04 -0.01 -0.01 +0.02 +0.01 -0.01 1.07 1.05 v' . 4.“, '1. a. 4’ W-.. ._....W—.._.._—...—.» fl‘ohvy . ‘. - ‘5'. v.'. Table 9. LiBr LiBr (con't.) 0.263 0.1315 0.079 0.058 0.026 0.68 0.511 0.341 0.171 0.102 0.068 0.034 0.419 0.314 0.210 0.105 0.063 0.042 0.021 0.440 0.330 0.220 0.110 0.066 76 Tetrahydrofuran 28 27 27 23 25 25 24 - 5o -53 -50 -53 - 52 0.55 0.57 0.55 0.55 0.59 0.62 0.60 0.64 0.62 0.69 0.59 0.64 0.47 0.45 0.45 0.38 0.42 0.42 0.40 -0.84 -0.89 -0.90 -o.89 -0.87 1.03 1.05 1.03 1.03 1.07 1.10 1.08 1.12 1.10 1.12 1.07 1.12 0.80 0.78 0.78 0.72 0.75 0.75 0.73 -0.51 -0.56 —0.57 -0.56 -0.54 4—7‘ _.._. w—q’.‘ H... a—F- “a: -. ,. ‘,A_f“' . A4,, 5 9 77 Table 9. (con't.) 0.044 - 51 -0.855 —0.52 0.022 - 50 -0.84 -0.51 Pyridine L10104 0.470 -165 -2.77 -2.31 0.066 -170 -2.85 -2.39 0.033 -171 -2.86 -2.40 LiBr 0.250 -225 -3.77 -3.31 0.187 -226 -3.82 -3.36 0.125 233 -3 90 -3 44 0 075 -228 -3 82 -3.36 0.050 -228 -3.82 -3 36 0.025 -230 -3.85 -3.39 0.0125 -235 -3 94 -3.48 2-Methylpyridine 110104 0.305 -125 -2.10 -1.65 LiBr <0.26 -185 -3.10 -2.65 an} .A- not . ' . ~. g —-.g 4 5 -- 4qu .uonninr .. Figure 8. 78 Lithium-7 chemical shifts of LiBr and LiClOl$ in various solvents: A. B. C. D. E. F. G. H. I. J. K. L. LiClOu in acetonitrile LiBr in acetonitrile LiBr in dimethylsulfoxide LiClOu in dimethylsulfoxide LiClOu in tetrahydrofuran LiBr in tetrahydrofuran LiClOn in acetic acid LiBr in acetic acid LiClOu in acetone LiBr in acetone LiClOu in pyridine LiBr in pyridine 8 (P PM) 79 3.0 A - of 40" ° o-o-O———'"" 2.0 "‘ \O\o\__o O B o C LO L-O BEIO‘D—o—J' o D '35- a 000-0 °—'—°-—E°- 000—0 0 60 0,0 Loco—o—-o 0 O- F 00-0-0—0 0 ° -I.O - | o 4—o————o--—— 0’07 J oc>o--o ° ° °"' -2.0 - K «re 0- -3.0 " L _o 0.0 . 0 01 . 1 1 i O 0.2 0.4 0.6 CONC. (MIL) ’ 'l‘. y l- “.144... ":dfi' “ ._ .u - 80 bromide solutions which dissociate into solvent-separated or free solvated ions at lower salt concentrations. In contrast to acetonitrile, the lithium-7 NMR in dimethyl- sulfoxide solutions shows essentially no difference between the lithium bromide and perchlorate, which indicates that the lithium ion is strongly solvated by this solvent. These NMR results correlate very well with the infrared solvation studies which indicate little effect of the anion on the inner solvation shell of lithium ion in dimethylsulfoxide solutions.5 The lithium-7 chemical shifts for lithium perchlorate and bromide solutions in tetrahydrofuran solutions are con- centration independent with the chemical shift of the bromide solution about 1.2 ppm downfield from the perchlorate solu- tion. This strong lithium-7 chemical shift dependence on the anion indicates strong lithium ion-bromide interaction. The strong dependence of the anion on the frequency of the lithium ion solvation band in tetrahydrofuran“ also indicates that the anion is associated with the lithium ion. It is known that ion aggregation and ion pairing occurs extensively in tetrahydrofuran.25'26 In acetic acid, also a solvent of low polarity, very little difference is observed between the lithium bromide and perchlorate solutions by lithium-7 NMR. In this case, the lithium ion must be largely solvated by the acetic acid to form free solvated lithium ions or solvent-separated ion pairs. “9.. .Q‘w-.-.m Ll) amt: ‘fl '\ .-&. 9‘ -jJ ' 4‘ 81 Studies of lithium perchlorate in acetone give somewhat different results in that the lithium-7 chemical shift of tzhis salt decreases at concentrations below 0.15 molar. At the lowest concentration studied (0.052) the chemical shift appears to approach to the concentration independent lithium-7 chemical shift in lithium bromide solutions. Pyridine solution results show a linear lithium-7 chemical shift dependence on the lithium salt concentration with about a 1 ppm separation between the lithium perchlorate and bromide salts. In the concentration range studied (0.6 to 0.2 molar), it would appear that there is little change in the relative concentrations of the species giving rise to the anion dependence of the chemical shift. Data obtained at 23.3 MHz for lithium salt solutions in several solvents are given in Table ‘10. These results show that 4-methylpyridine solutions of lithium salts show the expected behavior similar to that of pyridine. It should be noted that the chloride solutions in 4-methylpyridine do not give as great a downfield shift as do the bromides. The anion effect on the lithium-7 chemical shift in pyridine and 4-methylpyridine indicates that the same species may be involved in both the lithium-7 NMR and the infrared sol- vation band anion dependencies. The correlation of the sodium-23 ion chemical shift in 89 various solvents with solvent donor ability suggests that the lithium-7 chemical shift might also show similar results. However, this is not the case. As shown in Figure 9, the Jithimn-7’ion chemical shift in several solvents do not show 5'”... . 34:35.1 __ d“. ".101.- .2! 82 Table 10. Chemical shifts for lithium salt solutions in several solvents at 23.3 MHz. 0 m - gz -1x10'60bx .1x10“6 Pyridine LiClOu 0.1513 47.2 2.02 2.25 0.267 46.4 2.00 2.23 0.383 46.2 1.98 2.21 0.412 45.1 1.93 2.16 P ridine LiBr 0.140 67.2 2.88 3.11 0.212 67.2 2.88 3.11 0.269 67.4 2.89 3.12 0.375 67.8 2.91 3 14 <0.50 68.2 2.925 3.15 Pyridine LiCl 0.140 60.6_ 2.60 2.83 0.257 59.35 2.55 2.78 0.327 59.8 2.565 2.79 0.493 59.1 2.53 2.76 4-Methylpyridine LiClOu 0.175 47.65 2.04 2.26 0.378 47.1 2.02 2.24 0.532 46.6 1.98 2.20 -------_-__-9;f?f ....... ff:39-----_--3;??& ......... 3:32 ..... 4-Meth lpyridine LiBr 0.104 66.0 2.83 3.05 0.200 69.1 2.96 3.18 83 Table 10. (con't.) 0.424 68.55 2.94 3.16 0.689 68.3 2.93 3.15 4-Methylpyridine LiCl 0.234 58.0 2.49 2.71 <0.637 56.4 2.42 2.64 4-Methylpyridine LiNo3 0.469 43.9 1.88 2.11 2-Chloropyridine LiClOu 0.203 21.7 0.93 1.00 0.537 18.0 0.77 0.84 2-Chloropyridine LiBr <0.326 49.0 2.10 2.17 Dimethylsulfoxide LiClOu 0.383 -30.1 -1.29 -l.05 0.719 -30.0 -1.29 -l.05 Dimethylsulfoxide LiBr 0.0328 -3002 '1029 -1005 ' 9- tn "'1; ' . . z 5 - n h “alum \4 Q‘- 1. c t .' ‘- ' L Figure 9. 84 Lithium-7 chemical shift versus solvent donor number. A. B. C. D. E. G. Nitromethane Acetonitrile Acetone Tetrahydrofuran Dimethylsulfoxide Water Pyridine 85 n L. . In" I A. . skull... film-{kt}: : phi mmmZDZ 10200 on ON 0. q u q 8 d (de) 86 95 any correlation with Gutmann's donor numbers. Maciel, 23 a1. suggested that Kosower's Z valueglnian emprical solvent pol- arity scale based on the position of the charge-transfer bands of l-alkylpyridium iodide complexes in various sol- vents, gave a linear relationship when plotted against the lithium-7 chemical shift. Again the fit is not good for all cases for which data are available as is shown in Figure 10. There may be, however, some correlation of the Z values with the aprotic solvents pyridine, acetone, dimethyl- sulfoxide and acetonitrile. To study the lithium-7 chemical shift dependence on the anion, tetrabutylammonium bromide was added to 0.26 molar lithium perchlorate solutions in pyridine. The lithium- ? chemical shift is seen in Figure 11 to change smoothly upon addition of the bromide ion. The limiting shift cor- responding to that of lithium bromide solutions is reached above a ratio of bromide ion to lithium ion of 2.0. This insensitivity of the chemical shift to the excess bromide ion seems to indicate that one kind of cation-anion inter- action such as ion pairing is involved with no further for- mation of species such as bromide-lithium ion-bromide triple ions. As a check on the effect of tetrabutylammonium and per- chlorate ion on the lithium-7 chemical shift in these solu- tions, a second experiment was performed in which both the bromide and perchlorate concentrations were varied while the lithium and tetrabutylammonium concentrations were held at 0J24 and 0.47 molar respectively. The lithium-7 chemical ifl:~‘- ' I 4.&__;:. I,” ‘ .‘ID .‘ .1: f‘.’ O!“ ‘- 1‘3 .1» Figure 10. 87 Lithium-7 chemical shift versus solvent Z value. 10 20 30 L: 5. 6 7 Pyridine Acetone Dimethylsulfoxide ‘Acetonitrile Acetic acid Methanol Water (PPM) 88 D '/ _. c a E u— / . C r .. A / 1 l I l 50 60 70 80 90 Z VALUE IOO 89 shifts of these solutions are also shown in Figure 11, and are essentially identical to the results for the lithium perchlorate solutions to which tetrabutylammonium bromide was added. These results indicate that the lithium-7 chem- ical shift is largely governed by the ion pair formation with the bromide ion and is not affected by the presence of the perchlorate or tetrabutylammonium ions. l One of the experimental difficulties with lithium-7 NMR is the strong effect of small amounts of water on the chemical we”. ---w*.. .11.} shift. Figure 12 shows the lithium-7 chemical shift in w“- . . . water acetonitrile mixtures as a function of solvent compo- sition. The midpoint of the chemical shift transition occurs at about 0.01 mole fraction water in acetonitrile, which indicates very strong preferential solvation of lithium ion by water. This sensitivity of the lithium-7 chemical shift to small amounts of water in non-aqueous solvents requires extreme care to exclude water in the preparation of these samples. Consequently the nonaqueous solutions used in this study were prepared in a dry box. The dimethylsulfoxide-pyridine mixed solvent system was also investigated. The lithium-7 chemical shift of 0.38 molar lithium perchlorate in dimethylsulfoxide-pyridine mixtures is shown in Figure 13. It is seen that the chemical shift midpoint between the two pure solvents occurs at about 0.11 mole fraction dimethylsulfoxide, which indicates prefer- ential solvation of lithium ion by dimethylsulfoxide. These results are quite similar to those reported earlier94 for sodium-23 chemical shifts in the same solvent mixtures. @ Figure 11. 9o . The change in the lithium-7 chemical shift of lithium perchlorate solutions in pyridine with addadtetrabutylammonium bromide. D 0.26 pg LiClOu + BuuNBr 00.24 g Li+, 0.47 M BuuN+, Br- + C104- = 0.71. 8 (H2) 91 :mfit"'rt=*'“‘”- 1'- 1- 'm l 1 1 ___|__«r 1 0.5 LO L5 2.0 5.0 MOLE RATIO (BF/u“) 92 Figure 12. Lithium-7 chemical shifts in water- acetonitrile mixtures at 60.2 MHz. a P14?! ,. ...........: £151. in; on... 223...... 0.6: 6.. To 0.6 ...o «.o ..o o f 4 A . a O 0/ on! . 96 w .. o.. a. III III ' (mu) 9 94 Figure 13. Lithium-7 chemical shifts in dimethyl- sulfoxide-Pyridine mixtures. 23.3 MHz, 0.38 1_v1_ LiClOu 95 30*- 20- IO- ll'- 'D'U- '.'-'- .‘U'.'"'ll-"" I- -' O/ '50 L 0 -IO _ ~20 — -30 -40 «N1. a 04 00 DMSO FRACTION dz MOLE a- ‘. , .3.‘ O 96 Since the donor numbers of dimethylsulfoxide and pyridine are quite close, 29 and 33 respectively, this perferential solva- tion of lithium and sodium ions by dimethylsulfoxide is somewhat surprising. It has been shown by Brilluion scat- tering techniques101 that pyridine disruptes the highly ordered structure of dimethylsulfoxide. Thus with small amounts of added pyridine the dimethylsulfoxide structure is broken up and the dimethylsulfoxide becomes a better co- ordinating agent. The effects of temperature on the chemical shifts of .Lithium perchlorate and bromide solutions in acetonitrile were dertermined over the range of +20 to -45° C. As shown in Figure 1!», the chemical shift difference between the two solutions derxreased slightly with a decrease in temperature which indi- caiaes that the bromide--lithium ion interaction becomes less annortant. Similar studies in acetone and pyridine gave corisrtant chemical shift differences down to temperatures of -6C) sand -40° C respectively. Some line broadening was <finsexrved as the temperature decreased. Pyridine and acetoni- triILea solutions gave maximum linewidths of approximately 2 Hz Just: loefore the solutions froze, after which the resonance was too broad to detect. 'The solvent tetramethylguanadine (TMG) was investigated and found to have uniquely different characteristics. SOIutions of lithium perchlorate and bromide in this solvent hafire (Sbserved chemical shifts of -0.1 and -0.45 ppm which are temperature independent. The linewidths are significantly broadened compared to those observed in other solvents. As _—-_........-_..~......~u A—nb- was - .. ‘1 p p . h: ..n _-. M .1‘ fl Figure 14. 97 Lithium-7 chemical shift of lithium perchlorate and bromide in acetonitrile versus temperature. 00.5 E LiClOu 00.4 M LiBr A GCIOL, - 681‘ 98 LI... :11! L07 mink 6.... cu. c on O 4 . . +00 \\\\\D \D n. o. \\\\\\\\ 10m V D g 1. If so, he will enter the octal equivalent in the switch register (SR) and depress continue. The computer then stores SR in TIMES and again halts. He then enters the desired starting address in SR, 26999 for a sample smooth or 26919 for a ref- erence smooth. 116 If the operator wants only one smooth, he may after the first halt, enter either sample or reference smooth start- ing address (26999 or 26919) into SR and depress reset and start. The program then calls up the data starting addresses 19999 or 14999 (ASTAD or BSTAD), adds M4, the number of points on one side of the central point in the smooth and deposits . the storage address of the first smoothed point at STORE. 3 Then the NDPS is stored at data count (DCNT) and program i control is tranferred to subroutine LOADT. é Subroutine LOADT is used to load the first N data points E )7. into the temporary location where N is the number of points in the smooth and N': 15. LOADT first stores ASTAD or BSTAD which is still in AC9 at the first data point (ADR). NDPS is loaded in A01, negated and used as the counter. A09 is loaded at ADR and this is stored with respect to A02 which contains the starting temporary location TMPAD. The sub- routine loops after incrementing ADR, A02 (TMPAD), and A01 (-NDPS) until A01 = 9 when the temporary storate is full and a skip results in a jump to subroutine SMOTH. Subroutine SMOTH sets up the smoothing process and clears the partial sum (PSUM). The temporary location count (TCNT) is set equal to M4 and the constant address count KCNT is set = K9ADR + M4. The center point is first weighed by loading it into A01 and K9 in A02 with the arithmetic being done by calling subroutine SUM with the return to the loop sub-subroutine. The latter adds the temporary data at P 1:TCNT with the sum in A01 and the appropriate constant 0 loaded indirect at KCNT into A02 ending with a call to SUM. 117 The loop is iterated with deincrementing of KCNT and TCNT until TCNT = 9 which causes a skip to subroutine STO. The weighting and normalization of the points is done in subroutine SUM. Because of problems of adding and storing the double precision products and since the product is easily set up for division by loading the denominator in A02 and calling DIV the normalization is carried out immediately after the multiplication. Due to the positive remainder, with negative numerators the high order part of the product containing the sign bit is stored in SAVE. It is recalled after the division and if negative, the remainder is negated before addition to REM. The normalized partial sum in A01 is then added to PSUM. It was observed that round off errors from neglecting the remainder in the division from SUM caused an error in the least significant digit. The remainder (A09) is updated by addition to REM. Subroutine STO first corrects PSUM for the cumulative remainder by loading REM into A01, clearing AC9, setting AC2 = l, and calling SUM. The final PSUM in A01 is then deposited indirect at STORE. The data count (DCNT) is deincremented and the program looped through ADDl until DCNT = 9. Upon a 9 result TIMES is deincremented and if = 9, TIMES is restored to l and END is called. If TIMES 9 9, a call is made in the escape check subroutine in the FTS program which returns control to the FTS program if the teletype escape key has been depressed. This offers a convenient mechanism to halt the program if a large number, .3. .- '.‘i." 5‘): F AW» _" 1" I‘T'.‘ Ww—L- 7.- :r' a .v . ’“~ in?» 118 e.g., 26999, has been inadvertently deposited in TIMES by subroutine START. Finally, SR program control is trans- ferred to the address in SR which, hopefully, is the appro- priate sample or reference smooth starting address (26999 or 26919). The END subroutine prints out the following message: 9-point smooth number of times sample smoothed = 9 number of times ref smoothed = 9 NDPS = 3779 where the numbers are called from NDPTS, SCNT, RCNT, and NDPS. They are printed in decimal form by the integer and text print subroutines gleaned from the FTS program. The END subroutine ends with a final escape check to return to FTS program control if desired by the operator before calling a halt. T7 " 3 4‘: A. 119 FIGURE Al PROGRAM SMOOTH no FIGLJRE A 2 SUBROUTINE START am on 22 -H II I on PRINT STATUS l READ 8R IAIPL ' REF RCN‘I’ '0' I 80 N1 4. I STORE I ”TAO‘P N4 81' OR! OAOTAD+U4 l J 121 Table A1 A listing of program smooth 26000 26001 26002 26003 26004 26005 26006 26007 26010 26011 26012 26013 26014 26015 26016 26017 26020 26021 2 6022 2 6023 2 6024 26025 1 I AFTER GOLAY AND SAVITIESKY ANAL. I .EOT 9 POINT DIGITAL SMOOTHING ROUTINE FOR CHEM. FTS-16 36.1627(l964) P.R. HANDY 2/28/72 1 LIST OF SUBROUTINES FROM FTS PROGRAM '900220 .DUSR MPY = 220 000221 .DUSR DIV = 221 000201 .DUSR IPRINT = 201 000222 .DUSR TPRINT = 222 000355 oDUSR EXCHK = 355 026000 .LOC 26000 020535 SAMPLE: LDA 03ASTAD 3 START HERE FOR SAMPLE SMOTH- 010541 182 SCNT 024527 LDA 10M4 3 +4 107000 ADD 0:! 044584 STA loSTORE 1 10005 (14005) 024525 LDA laNDPS 1 3 DATA PTS TO BE SMOTHD 044533 STA I’DCNT 3 DCNT = DPS 000435 JMP LOADT ‘020526 REF: LDA 0:BSTAD I REF SMOTH:START HERE 010532 ISZ RCNT 000770 JMP .-IO 102420 START: SUBE 0:0 1 CLEAR ACD 024526 LDA l:SCNT 044526 STA laRCNT 3 RCNT 8 SCNT 040524 STA 0:SCNT J SCNT = 0 004535 JSR STATS 3 PRINT STATUS 063077 HALT 3 ENTER NUMBER OF TIMES DATA TO BE :SMOOTHED INTO SWITCHES IF >1 064477 READS l 3 AND DEPRESS CONTINUE; IF =1 IENTER ADRES IN SRJRESET18TART 044523 STA laTIMES 063077 HALT .1 ENTER ADRES 0F SAMPLE (26000) 30R REF (26011) INTO S.R. AND CONT. 074477 READS 8 3 READ SR INTO A03 001400 JMP 0:3 Table A1 (cont) 26026 26027 26030 26031 26032 26033 26034 26035 26036 26037 26040 26041 26042 26043 26044 26045 26046 26047 26050 26051 26052 26053 26054 26055 26056 26057 26060 26061 26062 26063 26064 26065 26066 26067 26070 26071 26072 26073 26074 26075 024502 124400 125400 030457 151400 021000 041377 125404 000774 022500 042474 010476 010476 000512 040473 024463 124400 030441 022467 041000 010465 151400 125404 000773 000477 054467 006220 030415 040466 006221 034464 175112 100400 030460 133000 050456 024457 107000 044455 002451 ADD]: LOADT: SUM: .EOT LDA NEG INC LDA INC LDA STA INC JMP LDA STA 152 132 JMP STA LDA NEG LDA LDA STA ISZ INC INC JMP JMP STA JSR LDA STA JSR LDA MUVL# 3:3 SEC NEG LDA ADD STA LDA ADD STA JMP 122 1:NOPTS 1:1 1:1 2:TMPAD 2:2 0:0:2 0:“1:2 1:1 53R 0'4 0:0ADR 0:0PADR ADR STORE SNOTH 0:ADR 1:NOPTS 1:1 2:1MPAD 0:9ADR 0:0:2 ADR 2:2 1:1 52R 0'5 SMOTH 3:RTN @MPY 2:NORM 0:3AVE QDIV 3:3AVE 0:0 2:PSUM 1:2 5050505. 505050“ 3 LOAD 9 '9 ‘8 TEMP ADRESS LOAD ACO AT AC2 STORE ACO AT A02 -1 TMPAD +1 +1 0: SKIP LOAD AT NEW ADR 3STORE LAST POINT 3 3 ADR +1 STORE +1 3START ADRESS bhhhhhhh be 3 LOAD 9 -9 ADRESS 0F TEMP LOAD DATA POINT STORE ACO INDIRECT AT A02 ADR = ADR +1 TMPAD +1 -NOPTS +1.5KP 0N R=0 SUBROUTINE SUM MULT K TIMES NUMBER LOAD NORM FACTOR RECALL HI ORDER PRODUCT JSKIP IF POSITIVE i 3 3 NEGATE REM IF NEGATIVE ‘LOAD PSUM ADD NEW TO OLD PSUM 2:PSUM 3 STORE NEW TOTAL PSUM 1:REM 0:1 1:REM GRTN 3 UPDATE REMAINDER Table A1 (cont) 26076 26077 26100 26101 26102 26103 26104 26105 26106 26107 26110 26130 26131 26132 26133 26134 961 35 26136 26137 26140 26141 26142 26143 26144 26145 26146 26147 26150 26151 26152 26153 26154 000012 000347 026100 000073 000066 000047 000016 177753 000000 000000 000000 000010 026111 000012 000017 000011 000004 000011 000004 003710 000010 026115 026121 .RDX 10 KOADR: .RDX 8 TMPAD: .RDX 10 TEMP: POINTS 123 NORM : 231 .+1 ' K0 : 59 K1 : 54 K2 : 39 K3 : 14 K4 : -21 KS : 0 K6 : 0 K7 : 0 0"! .BLK 15 9 CONST =POINTS/2 010000, 014000 000000 000000 000000 000000 000000 000000 000001 000000 000000 000000 000000 000000 000000 000457 INORMALIEATION FACTOR 3 CNTR PT NT'NG FACTOR ’ cp. +! u‘ u 3 CF: +2 3 CPO +3 3 CPC +4 3'21 3ASSIGN TEMP. LOCATIONS 39 POINTS IN SMOOTH 1 TEMP LOCATION 4 FOR 9-PTSMOTH FOR 2K WORDS CENTER SMOTHING POINT LAST ADR POINT OF TEMP SAMPLE START ADRESS REF START ADRESS CURRENT ADDRESS OF LAST DATA ADRESS TO STORE SMOTHD DATA -# PTS SMOTHED # TIMES SAMPLE SMOTHED REF SMOTHED COUNT # TIMES SMOOTH TO BE REPEAT RETURN PARTIAL SUM OF PRODUCTS HIGH ORDER PRODUCT REMAINDER 3CONSTANT COUNT TEMP WORD STORAGE NOPTS POINTS M4 :CONST 1 = NOPTS-1/2 = NDPS : 2001-POINTS 1 .RDX 8 P0: TEMP+CONST 1 PADR: TEMP+POINTS~1 1 ASTAD : 10000~ 1 BSTAD : 14000 1 ADR : 0 1 STORE : 0 1 DCNT : 0 1 SCNT : 0 1 KCNT: 0 1 0 TIMES TCNT: 0 TIMES: 1 1 kTN: 0 1 PSUM: 0 1 SAVE: 0 1 REM: 0 1 KCNT: ‘0 HOLD: 0 1 STATS: JMP END+2 .EOT Table A1 (cont) 26155 26156 26157 26160 26161 26162 26163 26164 26165 26166 26167 26170 26171 26172 26173 26174 26175 26176 26177 26200 26201 26202 26203 26204 26205 26206 26207 26210 26211 26212 26213 26214 26215 26216 26217 26220 26221 26222 26223 .26224 26225 26226 26227 26230 102420 SMOTH: 040771 040772 020751 040763 026751 030714 143000 040765 032711 004670 030743 020753 113000 025000 030737 112400 031000 050754 032752 004656 024751 032747 000401 004652 “14744 014735 000760 024740 102420 111400 004643 052723 014723 000607 014725 000405 102420 101400 040721 000404 006355 074477 001400 STO: OEOT SUBE 0:0 STA STA LDA STA LDA LDA ADD STA LDA JSR LOOP: LDA ADD LDA LDA SUB LDA STA LDA JSR LDA LDA JMP JSR DR? DSZ JMP LDA 0:PSUM 0:REM 0:M4 0:TCNT 1:9P0 2:KOADR 2:0 3 KOADR 0:KCNT 2:9KOADR SUM 0:TCNT 0:2 1:0:2 2:?0 0:2 2:0:2 2:HOLD 2:9KCNT SUM 1:HOLD 2:9KCNT 0+1 SUM KCNT TCNT LOOP 1:REM SUBZ 0:0 INC JSR STA DSZ JMP DSZ JMP SUBZ INC STA JMP JSR 0:2 SUM 2:05TORE DCNT ADDI TIMES 0+5 0:0 0:0 0:TIMES END GEXCHK READS 3 JMP 0:3 LDA 2:P0 PSUM = 0 REM = 0 50h TCNT = 4 LOAD CENTER POINT ADRES OF K0 M4 1 KONSTANT ADRESS KCNT = KOADR+4 LOAD K0 ' ~.u +~.u!u 3 AC2= CENTER PT ADRESS 3 +4 3 AC2 = CP +4 3LOAD AC1 9 AC2: 1=CP+TCNT CP -TCNT AC2 LOADED AT AC2 =CP-TCNT HOLD C.P. - TCNT LOAD K CALL SUM RECALL HOLD KOADR + 4 50605050805050 3 xc~1= KCNT -| 3TCNT=TCNT-1: SKIP ON 0 3 LOAD REMAINDER 3 CLEAR ACO 3MULT *1 3 NORMALIZE REM SUM 3ADD TO PSUM 3STORE SMOTHED POINT 3 ENOUGH POINTS? 3 NO:ADR +1 AND SMOTH 3 TIMES = TIMES '1 SUM 3 +1 3 TIMES RESTORED TO 1 3RETURN TO EXECUTIVE? 3 READ S.R. 3 JUMP TO START ADRESS C(331 1’ 61:32 26233 26234 26235 26236 26237 26240 26241 26242 26243 26244 26245 26246 26247 26250 26251 26252 26253 26254 26255 26256 26257 2 62 60 2 6261 26262 2 6263 26264 2?6265 26266 2 6267 2 6270 26271 2 6272 2 6273 26274 26275 26276 2 6277 26300 2 6301 2 6302 226303 004402 063077 054713 030475 0116272 020672 006231 030423 00622 020701 C06201 030440 006222 020674 006201 030453 006222 020660 006201 030455 006222 006355 002 667 024261 050055 044517 052116 051440 047515 052117 006510 052012 046511 051505 051040 043105 051440 047515 052117 042510 020104 020075 000000 ,-- Table A1 (cont) END: Ila—“‘0 JSR HALT STA LDA JSR LDA JSR LDA JSR LDA JSR LDA JSR LDA JSR LDA JSR LDA JSR LDA JSR JSR JMP 6+: .TXT 125 0+2 3:RTN 2:MSGA OTPEINT 11:N01"TS 91 PRINT 2: M561 OTPHINT 0:RCNT 911’1xINT 2:1‘1362 OTPKINT U: SCNT OIPHINT 2: M303 QTPRINT 0:'NDPS OIPRINT 2:MSG4 OTPRINT PEXCHK GRTN .-P 3 3 3 END OF SMOTH PROG CR & LF PRINT NOPTS LOAD ADRES OF MSG PRINT KCNT PRINT MSGZ PRINT SCNT PRINT M563 PRINT NDPS CR & LF IF ESC HAS BEEN DEPRESSED: 31:ILL RETURN TO EXECUTIVE. Table A1 (cont) 26304 26305 26306 26307 26310 26311 26312 26313 26314 26315 26316 26317 26320 26321 26322 26323 26324 26325 26326 26327 26330 26331 26332 26333 026305 005015 044524 042515 020123 040523 050115 042514 051440 047515 052117 042510 020104 020075 000000 026324 005015 042116 051520 036440 000040 026332 005015 000000 026013 M562: TI ME 5 5A MP LE 5 MO 01 HE M564: 0+1 oTXT 0+1 .TXT 0+1 oTXT .END 26013 126 '<15><|2> '<15><12> '<15><12> flUDl fiDR ASTAD BSTAD CONST DCNT END. HOLD K0 KOADR K1 K2 K3 K4 K5 K6 K7 KCNT LOADT LOOP M4 MSG] MSG? M563 M564 NDPS NOPTS NORM P0 PADR POINT PSUM RCNT REF REM RTN SAMPL SAVE SCNT SMOTH START STATS 5T0 STORE SUM TCNT TEMP TIMES TMPAD Table A1 (cont) 026026 026137 026135 026136 000004 026141 026231 026153 026100 026077 026101 026102 026103 026104 026105 026106 026107 026152 026044 026170 026131 026260 026312 026331 026337 026132 026130 026076 026133 026134 000011 026147 026143 026010 026151 026146 026000 026150 026142 026155 026013 026154 026211 026140 026057 026144 026111 026145 026110 127 "71171131111111Mimnmfiflmn“