SPECTROSCOPIC STUDfES 0F ALKAU AND' AMMONEUM IDN SOLVATION IN NON - AQUEOUS SGLVENTS Thesis for the Degree of Ph. D. MlCHIGAN STATE UNIVERSITY JOHN LYNN WUEPPER 1369 L I B R A R Y Michigan State University This is to certify that the thesis entitled SPECTROSCOPIC STUDIES OF ALKALI AND AF-iI'vIONIUM ICN SOLJATION IN NON-*AQUEC US SOLVENTS presented by John Lynn Wuepper has been accepted towards fulfillment of the requirements for PhoDo degree in Chemiitry Major profesy / / i. ' '1 . r Date 41.9“) *1 4; ‘/ 0-169 ABSTRACT SPECTROSCOPIC STUDIES OF ALKALI AND AMMONIUM ION SOLVATION IN NON—AQUEOUS SOLVENTS By John Lynn Wuepper The subject of this thesis is the study of alkali and ammonium salt solvation in non-aqueous solvents by infrared and nuclear magnetic resonance‘spectroscopies. In 2-pyrrolidones, solutions of a given cation show anion independent far infrared bands which have been ascribed to the vibration of the cation in a solvent cage. For example, in l-methyl—2-pyrrolidone (1M2PY) solutions, the bands occur at 398, 2ou, 207, 140, and 106 cm"1 for lithium, sodium, ammonium, potassium and rubidium ions respectively. The validity of the above assumption was tested by isotopic substitution. For example, the L17 band falls at 398 cm-1. In changing the cation to L16, the frequency shifts to 420 cm_l. Similarly, the band frequency shifts from 207 cm‘1 to 191 cm"1 when ND“+ is substituted for NHu+. Examination of the lM2PY spectrum in the sodium chloride region showed that the position of the carbonyl band is shifted to lower frequencies in the presence of John Lynn Wuepper alkali metal salts and, therefore, it is evident that the coordination occurs through the carbonyl oxygen. In an attempt to determine alkali ion coordination numbers in lM2PY and dimethylsulfoxide (DMSO), mole ratio studies were carried out in dioxane, a relatively inert solvent. Either the infrared band frequency or the nmr proton shift was followed as the ligand/cation ratio was varied. The results showed a coordination number of 4 for lithium in lM2PY, 4 for sodium in lM2PY and 6 for sodium in DMSO. The absence of Raman lines in lM2PY solutions of lithium salts suggests that the ion-solvent bond is ionic. The strength of the ion-solvent bond is illustrated by the fact that solid solvates of the alkali and ammon- ium salts with lM2PY can be prepared. SPECTROSCOPIC STUDIES OF ALKALI AND AMMONIUM ION SOLVATION IN NON-AQUEOUS SOLVENTS By John Lynn Wuepper A THESIS Submitted to Michigan State University in partial fulfillment of the requirements for the degree of DOCTOR OF PHILOSOPHY Department of Chemistry 1969 ACKNOWLEDGMENTS The author would like to sincerely thank Dr. A. I. Popov for the interest and enthusiasm which he offered during the course of this work. Appreciation is extended to Dr. M. C. Day, Dr. G. E. LeRoi, Mr. J. F. Lyndrup, Dr. B. W. Maxey, and Dr. P. G. Sears for their technical aid and constructive suggestions. The author would like to express his gratitude to Michigan State University, the Socony Mobil Oil Corpora- tion and the Dow Chemical Company for financial support. ii TABLE OF CONTENTS Page ACKNOWLEDGMENTS. 11 LIST OF TABLES . . . . . . . . . . . . V’ LIST OF FIGURES. . . . . . . . . . . . vi LIST OF APPENDICES. . . . . . . . . .viii Chapter I. INTRODUCTION . . . . . . . . . . 1 II. HISTORICAL PART . . . . . . . . 3 Far Infrared Studies of Ionic Solvation 3 Nuclear Magnetic Resonance Studies of Ionic Solvation . . . A Special Problems in the Determination of Solvation Numbers . . . . 5 Pyrrolidones . -. . . . . . 8 Far Infrared Spectroscopy.‘ . 9 III. EXPERIMENTAL PART . . . . . . . . . 12 Far Infrared Measurements. . . . 12 Infrared Measurements, A000- 625 cm-1 . . 15 Ultraviolet Spectra. . . . . . . . 15 Laser Raman Spectra. . . . . 16 Nuclear Magnetic Resonance Measurements . 16 Gas Chromatography . . . . . . . . 1? Purification of Solvents . . . . . . 18 Preparation of Salt Solutions . . . . 21 Salts . . 21 Preparation of Alkali Metal- lM2PY Solid Solvates . . . . . . . . . . 26 Analyses . . . . . . . . . . . 27 IV. RESULTS AND DISCUSSION. . . . . . . . 28 Infrared Solvation Studies of Alkali and Ammonium Ions in 2-Pyrrolidones. . . 28 iii Chapter Correlation of Ion-Solvent Band Fre- quencies with Physical Properties of the Ions . . . . Infrared Anion Studies. Study of the Carbonyl Frequency in. lM2PY— Dioxane Mixtures. . Solid Solvates of Alkali Salts with lM2PY . . Raman and Ultraviolet Spectra of Alkali Ion Solutions. . Nuclear Magnetic Resonance Study of Alkali Ion Solvation . . Far Infrared Study of NaAl(But)u in Various Solvents. V. MISCELLANEOUS OBSERVATIONS 1- -Methyl— 2- Pyrrolidone Complex with TiFu. Alkali Salt Solvation Studies in Other Solvents . Low Temperature Spectra of Acetone Solutions of Sodium Iodide and Ammonium Thiocyanate . VI. SUGGESTIONS FOR FURTHER STUDY BIBLIOGRAPHY. . . . APPENDICES . . . . . . . . . iv Page 52 5A 57 63 66 67 95 99 99 99 102 104 107 113 Table 10. ll. l2. 13. l“. 15. LIST OF TABLES Page Cation coordination numbers obtained by integration of coordinated solvent proton resonance. . . . . . . . . . . . 6 Solvation numbers of alkali ions by proton NMR. . . . . . . . 6 Hydration numbers of sodium ion obtained by different methods . . . . . 7 Physical constants of 2—pyrrolidones. . . 9 Absorption spectra of lithium ions in 2-pyrrolidones . . . . . . . . 32 Absorption spectra of alkali metal ions and ammonium ions in 2-pyrrolidones . . . . . 33 Beer's law data for LiClOu and NaClOu in lM2PY . . . . . . . . . . . . A2 Absorption spectra of Li6 salts in lM2PY . . A3 Infrared absorption data: anion study . . . 58 Solid complexes of lM2PY with alkali and ammonium salts . . . . . . . 65 Coordination number of Li ion in lM2PY . . . 76 Infrared spectra of NaA1(But)u in various solvents from 320 cm-1 to 170 cm-1 . . . . 98 Absorption spectra of alkali metal and ammonium ions in propylene carbonate and tri-n-butylphosphate . . . . . . . . . 102 Solvents which were eliminated for far infrared studies . . . . . . . . . . 103 116 Program cal . . . . . . . . . . . Figure LIST OF FIGURES Far infrared spectra. A. Pure l-methyl- 2-pyrrolidone. B. 0.3 M LiClOu in l-methyl- -2-pyrrolidone . . Far infrared absorption bands of 0.5 M alkali metal ions in l-methyl- -2-pyrrolidone Absorbance Kg. concentration for LiClOu in 1M2PY. Path length = 0.05 mm. Absorbance readings taken at 398 cm-1. . . . . . Frequency of the solvent-cation absorption band lg, the lM2PY/LiClOA mole ratio in dioxane solutions. LiClOu concentration = 0.5 M- . . . . . . . . m +111 1 ///cation solvent YE: wavenumber mcation msolvent of infrared band maxima for alkali and ammonium ions in 1M2PY and DMSO . . The carbonyl band of 1M2PY at various ratios of lM2PY/Lil. Concentration of 1M2PY = 2 M. Lithium iodide concentration variable Chemical shift of the N—methyl protons of lM2PY vs. lM2PY/LiClOu mole ratio. Concen- tration of LiClOu varied. Concentration of lM2PY = 2 M. . . . . . Chemical shift of the N-methyl protons of lM2PY vs. lM2PY/LiClOu mole ratio. LiClOu concentration = 0.5 M. Concentrations of 1M2PY varied. A. lM2PY + LiClOu solutions in dioxane. B. 1M2PY in dioxane . Chemical shift of the N-methyl protons of lM2PY Kg. lM2PY/Lil mole ratio. Concentra- tion of LiI varied. Concentration of lM2PY=2M. . . . . . . .. vi Page 29 35 40 “9 55 60 70 72 7A Figure 10. ll. 12. 13. 14. 15. l6. 17. 18. Page Chemical shift of the dioxane protons 1g. lM2PY/LiClOu mole ratio. 1M2PY concentra— tion = 2 M . . . . . . . . 77 A. Chemical shift of DMSO protons vs. DMSO/NaAl(But)u mole ratio, NaAI(BuEIu con- centration = 0.36 M_. B.. Blank . . . . . 81 Chemical shift of DMSO protons vs. DMSO/ 84 NaAl(But)u with the baseline subtracted. A. Chemical shift of the N-methyl protons of 1M2PY vs. lM2PY/NaAl(But)u mole ratio. NaAl(But)u concentration = 0.38 M. B. Blank. 86 A. Chemical shift of dioxane protons vs. lM2PY/NaAl(But)u. NaAl(But)4 concentration = 0.36 M. B. Blank. . . . . . . . . . 89 A. Chemical shift of N-Methyl protons of 1M2PY Kg. lM2PY/(Hex)uNBr mole ratio. Salt concentration = 0.38 M. B. Chemical shift of terminal methyl protons of NaA1(But)u vs. lM2PY/NaAl(But)u mole ratio . 91 NMR spectra of terminal methyl protons and methylene protons not bonded to A1 of tetrabutylaluminate anion in: A. DMSO/ NaA1(But)u of 0.25/1. B. DMSO/NaAl(But)4 of 18/1. Salt concentration = 0.36 M . . . 93 Far infrared spectra: A. DMSO/NaA1(But)u ratio of 5.5:l. B. DMSO/NaAl(But)u ratio of 18. C. NaA1(But)u in dioxane. D. NaAl(But)u in n- heptane. Salt concentra- tions = 0.36M. . . . . . . . 96 Sample data cards for program cal. 118 vii LIST OF APPENDICES Appendix Page I. Computer Program for Spectrophotometer Calibration. . . . . . . . . . . 115 II. Sample Frequency Calculations. . . . . 120 viii I. INTRODUCTION The behavior of electrolyte solutions has been the subject of numerous investigations, and a wide variety of techniques have been employed to study ion-solvent inter- actions. It has recently been pointed out that far infrared and nuclear magnetic resonance spectroscOpies offer a new approach to the study of ionic solvation (1-7). In choosing a class of solvents with which to work, two simple conditions had to be met. The solvent had to have useful transparency in the infrared from about 650 cm"1 to 50 cm-l, and the solvent had to be capable of dissolving a wide variety of salts. Preliminary investi- gations showed that the 2-pyrrolidones met these require- ments. The 2-pyrrolidones are good dissociating solvents and have been widely used in other studies such as electri- cal conductance (8) and reaction kinetics (9, 10). In the study of ionic solvation, we chose to start with the alkali metal ions so as to keep the system as simple as possible and to avoid formation of coordination compounds. The alkali ion-pyrrolidone system lends it— self to the study of ionic solvation and, therefore, was used for most of the research to be described in this thesis. Ammonium salts were also included in the study. Iull Illllnllll ll‘lllll The purpose of this work was simply to use the above and similar systems to study the phenomenon of ionic solvation in non-aqueous solvents. F‘ mar-x ~. II. HISTORICAL PART Far Infrared Studies of Ionic Solvation The history of far infrared spectroscopic observa— tions of interionic, intermolecular and ion-solvent vibra- tions is relatively brief. Evans and Lo (11) observed bands in the 100 cm.1 region for tetralkylammonium halides in benzene solution. Band maxima were observed at 120 and 80 cm"1 for (n-CuH9)uNCl and (n-CHH9)NBr respectively. Frequency shifts with the change in mass could be cor- related with the model of vibration of an ion-pair or of higher ionic aggregates. The study was not extended to polar solvents or other non-polar solvents. Edgell and co-workers (A—6) studied infrared bands of lithium, sodium and potassium tetracarbonylcobaltates in tetrahydrofuran solutions. The respective frequencies were A07, 190 and 150 cm-1. Frequency shifts were ob- served when the tetracarbonylcobaltate anion was replaced by Mn(CO)5’. The authors postulated that the bands were due to the cations vibrating in a cage consisting of solvent molecules egg the corresponding cation. Maxey and Popov (1-3) investigated anion independent alkali and ammonium ion bands in dialkylsulfoxides. For +, NHu+, Na+, K+, Rb+, and Cs+ in dimethylsulfoxide 3 Li the maxima occurred at H30 cm"l, 214 cm— 1 l l l , 200 cm” , 15A cm“ , 123 cm‘ and 110 cm‘l, respectively. Small shifts from these values were observed in the other dialkylsulfoxides. In this case, the anion apparently did not enter into the solvent cage around the cation. French and Wood (7) reported that sodium tetra— phenylborate exhibits a band at 175 cm"1 in pyridine, l,A dioxane,piperidine and tetrahydrofuran. They sug- gested that the bands were due to the vibrations of con- 1 tact ion pairs, and that the cations were not solvated in the above solvents. It should be noted that they did not study the dependence of the band frequencies on the anions of the salts used in their investigation. Nuclear Magnetic Resonance Studies of Ionic Solvation In addition to infrared techniques, nuclear magnetic resonance spectroscopy has been proven to be very useful in the study of metal ions in solution. The review of Hinton and Amis (l2) thoroughly surveys the techniques used and the results obtained through 1966. One of the most significant advances in the chemistry of solutions is the determination of primary or inner sphere solvation numbers. The chemical shift of a proton is very sensitive to the electronic environment in which that proton finds itself. In a salt solution, different signals can often be observed between bulk solvent molecules and solvent molecules which are involved in solvating the salt. Coordination numbers are determined by direct integra— tion of the coordinated solvent resonance (13-18). The feasibility of this type of experiment is dependent on the rate of solvent exchange, and the strength of the ionic solvation. Low temperatures, on the order of 0° to —100°, are often employed to slow down the exchange rates so that two solvent resonances can be detected. Examples of ions which have been successfully studied by this technique are listed in Table l. The alkali metals are conspicuously absent from the list. Alkali ion solvation numbers have been measured by proton nmr. The measurements are based on either relaxation times as a function of salt concentration (19, 20) or on the chemical shift (2, 21, 22) caused by the alkali ion. The results for several alkali ions are listed in Table 2. The agreement between methods leaves much to be desired. The fact that only small solvation numbers are obtained with nmr indicates that inner sphere solvation numbers are being measured. Special Problems in the Determination of Solvation Numbers Historically, one of the main problems in the determination of solvation numbers has been that of differentiating between inner sphere coordination and 1" IIIIII.‘ I'll! ill. 'Il‘IIIT.I I I. l "l‘ I ll TABLE l.——Cation coordination numbers obtained by integra- tion of coordinated solvent prOton resonance. Cation Metal Ion Solvent Coordination Anion Reference Number Be(II) DMF A 0104— 13 Mg(II) Water u,5.7 0104’ IA, 15 Methanol CO(II) DMF 6 010”- 16 A1(III) DMSO 6 C10q_ l Al(III) Water* 6 0104— 18 n 6 Cl- 11 n O Br.“ n H 6 IJQ3-_ It Ga(III) Water 6 ClOL4 " In(III) Water 6 C104— " Be(II) Water A Clo”- ” *These values deviate from 6 slightly as the tem- perature is varied. TABLE 2.—--sOlvationa numbers of alkali ions by proton NMR. Solvation Numbers Cation Proton Relaxation Chemical Shift Li 1 (l9) 5 (20) u (21) 2 (DMSO) (2) Na 3.6 3 3.1 l,u (THF) (22) K 6b 1 2,1 Rb 9.9 — 1.6 Cs 19.6 — 1.0 aSolvation numbers in water except for references 2 and 22. b e . .. ' , V o Assumed as a bas18 for calculation. xi outer sphere coordination. Inner or first sphere coordi- nation refers to the solvent molecules bonded to the ion. All other solvent molecules involved in solvating the cation can be considered to be in the secondary solvation shell. The problem is well illustrated by examining the results of different experiments for the determination of sodium ion hydration numbers. The results are taken from the review on solvation by Bockris (23) and are contained in Table 3. The small values are a measure of the primary solvation, and the large values are a measure of the total solvation. TABLE 3.-—Hydration numbers of sodium ion* obtained by different methods. Method Hydration Number Mobility A Water transport 3 Mobility 71 Dialysis 1? Partial molar volume 1 Polarimetric 700 approx. *Ref. 23. Pyrrolidones Pyrrolidones form complexes with a wide variety of metals including the alkali metals (2A, 25). Other com- plexes include those with salts of nickel, chromium, calcium, strontium, aluminum, zinc, cadmium, indium, antimony and cobalt (25-27). Lactams are cyclic amides of the following general L structure (I). (2H0) \C/N‘R l-methyl—2upyrrolidone, 2—pyrrolidone and l—vinyl— 2—pyrrolidone are lactams with structures II, III, and IV respectively. Table A contains pertinent physical con- stants of the compounds. 2H CH2 2HC CH2 2H CH2 H _ HC N --H HC N—c =CH2 II II 0 O 0 II III IV l-methyl-2-pyrrolidone 2-pyrrolidone l-vinyl—Z-pyrrolidone (1M2PY) (2PY) (1V2PY) 9 TABLE A.--Physica1 constants of 2-pyrrolidones. Dielectric Dipole Molecular Meltin B 111 S 8 0 n8 °1vent Constant Moment Weight Point Point 1M2PY 32 (28) A.09(29) 99,13 _2uoc(31) 202(31) lV2PY 22.A(32) —_ 111.1 ”110 _- Dyke and co-workers (8) studied the conductance of sodium, potassium and some organic salts in l-methyl- 2-pyrrolidone at concentrations of approximately 10"3 M. The Fuoss-Onsager equation (33, 3A) was applied t0 the data and the ion pair formation constants were calculated to be zero for all salts studied. The results indicate that l-methyl—2-pyrrolidone is a very good dissociating solvent. Zaugg (10) studied the catalytic effect of l-methyl- 2—pyrrolidone on the alkylation of enolate anions and concluded that sodium ion strongly interacts with the N electron system of the 1M2PY molecule. Far Infrared Spectroscopy Good quality, double beam, far infrared absorption Spectrophotometers have been commercially available since about 1961. Far infrared work has always been hampered by the lack of a high intensity source. The two most practical 10 sources available now are the globar, for work above 100 cm'l, and the mercury lamp for work below 100 cm'l. Some real advances have been made in detector technology, such as the doped germanium detector, but expense, the necessity of liquid helium and limited frequency response still favor the Golay detector. The subject is well treated by Smith (35). L, Since the advent of the first high quality double beam instruments, advances in commercial units have been #4 limited to perfecting the original designs. This in itself is a significant development. For example, with some of the older instruments, water vapor present in the atmosphere had to be slowly purged from the spectro- photometer with a dry gas. This had to be done whenever optics changes were made. Purging time alone would amount to an hour at the very minimum if maximum source energy was to be obtained. With the new instruments, water vapor can be removed from the sample compartment as well as the entire instrument by vacuum evacuation (36). An entire spectrum from A00 to 30 cm.1 can be ob- tained in less than ten minutes. Just as important as the development of instru- mental design is the development of far infared refer- ence spectra (37-39). Compilations of far infared spectra have become available only recently. More 11 complete reference works will undoubtedly appear in the near future. E. D. Palik's (A0) far 1r bibliography covers the years 1892 to 1960. Stewart (Al) has written a very excellent general discussion on the subject. The re- view of Brasch e§_e;, (A2) surveys the literature pub- lished in 1965 and 1966. This review, as well as the recent article by Farraro (A3) discusses types of far 1r work and latest developments in the field. The alternative to far ir absorption spectroscopy is Fourier transform spectroscOpy. A good general history has been written by Lowenstein (AA). Better energy, resolution and speed are advantages over absorption spectroscopy. The major disadvantage is the computations involved in obtaining a spectrum from the Fourier transform. III. EXPERIMENTAL PART Far Infrared Measurements The far infrared measurements were performed with the Perkin Elmer 301 far infrared spectrophotometer (A5). I The range of the instrument is approximately 660 cm"1 to 13 cm’l. All spectra were obtained with the instrument f Operated in the double beam mode. Either purified tank nitrogen or liquid nitrogen, additionally dried by pas- sage through an Aquasorb drying tower, was used to purge the instrument of water vapor when measurements were taken below 330 cm-1 since water vapor absorbs strongly in this region. The gain of the spectrophotometer amplifier usually can be decreased about two steps after the instrument has been purged for several hours. It was found that liquid nitrogen is superior to tank nitro- gen in that lower amplifier gains are needed and better baselines are obtained. However, tank nitrogen is handy for running survey spectra. Demountable liquid cells, obtained from Barnes Engineering Co., Stamford, Conn., were found to be most useful when working with solvents with low vapor pres- sure. Teflon spacers ranging from 0.025 mm to 0.5 mm were used. The liquid cells can be easily filled with 12 13 a minimum of sample exposure to the atmosphere by using small volume syringes. Windows made of high density polyethylene are very practical since polyethylene is inexpensive and the windows can be frequently replaced. Polyethylene, how— ever, can be easily contaminated by certain solvents, and spurious absorption bands can result. The correct window size can easily be cut from a polyethylene sheet by using a KBr pellet press and a short section of sharpened high carbon steel pipe of proper diameter. If the solvent of interest is not completely transparent in a spectral region where an alkali ion- solvent band exists, the baseline must be subtracted from the spectrum. In order to avoid baseline subtrac— tion, an equal thickness of solvent can be employed in the reference beam provided no strong solvent bands are present in the region of interest. It is easier to construct matched cells with the Barnes demountable holders than to find matched molded polyethylene cells. The rigid 2 mm polyethylene windows can be used with measured teflon spacers. The thickness of the spacers can be checked with a micrometer. If difficulty in obtaining flat baselines is encountered, it often helps to simply dismantle and reassemble the cells.. 1A The path length was held constant by not dismantl- ing the cell between runs. Between measurements the cell was thoroughly flushed with acetone, pentane and dry air. It was then heated in an oven at 110° for thirty seconds, reflushed with dry air and filled for the next run. Reproducible absorbance measurements before and after cell cleaning were obtained in this way. A word of caution should be added about the use of matched cells. If a cell containing solvent is used to attenuate the reference beam, it is possible to observe negative absorption if the same solvent containing a dissolved sample in a matched cell is present in the sample beam. The reason is that with concentrated solu— tions there is less solvent in the sample beam cell. Higher gains may be necessary with a resultant increase in noise. Water vapor was used as the standard with which the frequency scale of the instrument was calibrated. A computer program which greatly aids in the presentation of the calibration data, as well as a more detailed dis- cussion of the calibration procedure for the 650 cm.1 to 13 cm'1 region, have been included in the appendix. A very convenient quick check calibration method has been developed recently (A6). A sample of yellow HgI2 is prepared in a polyethylene matrix. Yellow HgI has five well characterized bands in the 600 cm"1 to 2 15 50 cm-1 region. The preparation of this sample is very simple, the concentration of the standard is always the same and operating conditions do not have to be changed in order to check the calibration. Low temperature far infrared spectra can be ob— tained by special techniques. A Beckman low temperature cell with cesium iodide windows was used. A hole was cut in a spare sample compartment cover to accommodate the Dewar flask which fits on top of the cell. The cell it— self is contained inside a large glass tee with poly— ethylene windows which can be fully evacuated to prevent the condensation of moisture. The temperature is mea- sured by means of a thermocouple. Infrared Measurements, A000-625 cm-1 1to 625 cm-1 were Infrared measurements from A000 cm- obtained either with a Unicam Sp. 200 or a Perkin Elmer model 237 B spectrophotometer. Sodium chloride salt plates were used for both mull and solution spectra. Potassium bromide windows were also used for some solu- tion spectra. The slowest scan speeds available were used. A thin film of polystyrene was used for calibra- tion. Ultraviolet Spectra Ultraviolet spectra were recorded on a Cary 1A spectrophotometer. One centimeter cells were used. 16 Laser Raman Spectra Raman spectra were measured on a laser Raman spectrophotometer which was constructed at this uni- versity. The 51A5 A and A965 A exciting lines of a Spectra—Physics Model 1A0 Argon Ion laser were employed. Nuclear Magnetic Resonance Measurements Varian Associates A-60 and A56/60 instruments were used to record spectra. The A-60 was used in all the work involving l-methyl-2-pyrrolidone. The A56/60 was used in subsequent work in the dimethylsulfoxide solva- tion studies. Tetramethylsilane was used as an internal standard for all samples. Dioxane was used as solvent. A model 200 AB Hewlett Packard audio oscillator was used as an external audio frequency source for cali- bration purposes. A Heath Universal Digital Instrument counter, model number Eu-805 was used to monitor the audio frequency introduced. The advantage of the side band method is that it produces a reference sideband from TMS which can be placed any distance from a resonance of interest. This enables one to use the narrowest sweep width, 50 cps, and thus reduces the reading uncertainty of the chemical shift. For example, the dimethylsulfoxide protons occur at approximately 150 cps downfield from TMS. The 250 cps l7 sweep width must be used. On the 250 cps sweep width, one cps is equivalent to approximately 2 mm. If the ex— ternal sideband method is employed, the 50 cps sweep width can be used and one cps is equivalent to 10 mm. The decrease in relative reading uncertainty is obvious. Of course the sideband will be only as reliable as TMS itself. A counter gate time of ten seconds is desirable I, so that the frequency can be monitored to the nearest tenth of a cycle per second. The Heath counter did not 11 record a change larger than 0.1 cps over a several hour operating time. Gas Chromatography An F&M model 700 gas chromatograph equipped with a hydrogen flame detector was used to check solvent purity. A six foot 1/8 inch inside diameter column packed with 10% SE30 on Chromosorb S was used. The injection port temperature was maintained at approximately 10° above the boiling point of the solvent to be tested. Column tem- peratures were not kept at a constant value from run to run so that retention times were not recorded. Attenua- tion was greatly reduced on both sides of the major peak in order to be able to detect small amounts of impuri- ties. Nitrogen was used as a carrier gas, and the flow rate was calibrated with the aid of a bubble meter at 25 ml/min. l8 Purification of Solvents l—methyl—2-pyrrolidone, obtained from Aldrich Chemical Co., was stored over Linde AA molecular seives for several days before use. It was then decanted into a distillation pot, fresh molecular seives were added, and distillation was carried out through a 75 cm packed column at 1 mm and A8°. 'fi It was found that barium oxide decomposes 1M2PY ‘ under reflux conditions to produce a violet solution. Calcium sulfate is attacked by lM2PY and forms a white suspension under reflux conditions. All distillation glassware was flamed out before use to help minimize further moisture contamination. Silicon carbide boiling chips were found to be more efficient in preventing severe bumping than other more conventional boiling chips. The silicon carbide chips can be regenerated after use by heating to white heat in the flame of a Meker burner. Since bumping was a problem with this solvent, stirring with a magnetic stirrer was also employed. With the precautions described above, the final distillate was reduced to about 0.005 M water as deter- mined by Karl Fischer titration. Gas chromatography was used to check the purity of the distillate. Only one major peak was observed. Occasionally, a small peak which comprised less than 0.5% of the total area was 19 observed. It was not attempted to identify the impurity or determine its source. The solvent is hygroscopic and should be kept tightly capped under a dry nitrogen atmosphere. The sol— vent also decomposes at room temperature and should be used immediately after purification. The far infrared solvation bands are not affected by using unpurified lM2PY. However, any mole ratio re- sults involving 1M2PY would be dependent on the purity of the solvent. For example, in order to be able to calculate the lM2PY/salt ratio, the purity of both 1M2PY and the salt must be known. Yellowish crystals of 2-pyrrolidone, as obtained from the Aldrich Chemical Co., were purified by repeated fractional freezings until a colorless product was ob- tained which froze at 25.6° i 0.2° which is in good agreement with the literature value (30). The water content was 0.02 M_as determined by Karl Fischer titra- tion. When the solvent was being used it was stored in an oven at A0°. This temperature was sufficient to main- tain a liquid form long enough to obtain a far infrared spectrum. l-vinyl—2-pyrrolidone was obtained from the Aldrich Chemical Co. and used without further purification. Dimethylsulfoxide was obtained from Mathieson Cole- man & Bell Co., Cincinnati, Ohio. The solvent was 20 refluxed over calcium hydride for twenty four hours and then distilled at 1 mm pressure through a 75 cm packed column. Under these experimental conditions barium oxide decomposed dimethylsulfoxide to a black tarry mass. The distillation with calcium hydride proceeded smoothly to completion. Karl Fischer titration showed the solvent to be 0.006 M_in water. 1,A Dioxane was refluxed over metallic sodium for twenty four hours and distilled through a 75 cm packed column at one atmosphere pressure. The middle 70 per— cent of the batch was tested for purity by gas chroma— tography. Only one peak was observed. The solvent was stored over Linde AA molecular sieves, and the water con— tent was less than 0.005 M_as determined by Karl Fischer titration. Tri-n-butylphosphate was stored over calcium sule fate for several days before it was distilled at 1 mm of mercury. The water content of the purified material was on the order to 0.01 M, Barium oxide was found to de- compose tri-n-butylphosphate under reflux conditions since after several hours the liquid turned to a white Jelly-like mass. Propylene carbonate was refluxed over barium oxide for about twelve hours and distilled under vacuum at 59°. 21 Benzene, pentane and heptane were stored over sodium ribbon to remove water. Sodium metal can be disposed of by slowly adding it to isopropanol which has previously been dried over molecular sieves. It is important for safety reasons to use relatively dry isopropanol. Preparation of Salt Solutions For the mole ratio studies, the ligand/salt ratios were prepared on a weight basis. Solutions containing NaAl(But)u were prepared in a dry bag. salts The method of Schaschel and Day (A7) was used to prepare sodium tetrabutylaluminate. Tri-n-butylaluminum was obtained from K&K Laboratories in a one half pound lecture bottle. All the aluminum alkyls are highly reactive and tri-n-butylaluminum spontaneously reacts with moisture and oxygen in the air. On a weight basis, the heat given off from burning alkyl aluminum compounds is comparable to that of burning gasoline. If the com- pounds come in contact with the skin, severe burns are the immediate result. Information on the use, handling and chemistry of aluminum alkyls can be obtained from the Ethyl Corporation, Baton Rouge, Louisiana. All handling of tri-n—butylaluminum should be done in a dry box or dry bag filled with dried deoxygenated 22 nitrogen. The usual drying agents such as barium oxide may be used to dry the nitrogen. BAFS catalyst, obtain— able from Badische Analin Co., N.Y., N.Y., has been found to be an efficient deoxygenating agent. The catalyst is composed of copper impregnated on a pumice base. It can easily be regenerated by heating under a stream of carbon monoxide or hydrogen. The reaction for the preparation of NaA1(But)u is as follows: Heptane 3N8 + AAl(But)3 --*R§fIfii---- 3NaAl(But)u + Al It is important to carry out the reaction under a controlled atmosphere. All glassware was thoroughly flamed out to remove traces of water. The transfer of the tri-n-butylaluminum from the lecture bottle to the dropping funnel was made in a dry box. The preparation of the dispersion and the synthesis of the salt was carried out on a manifold system which provided a source of dried deoxygenated nitrogen and vacuum. A positive pressure of nitrogen was applied to the system by passing the gas over the reaction mixture. A very good sodium dispersion can be made by follow- ing the method described in "Inorganic Synthesis" (A8). Seventeen grams of metallic sodium were weighed into a 500 ml Morton flask and covered with n-heptane which had been dried over sodium wire. A Morton flask has 23 indentations in the side to aid the high speed stirring (A8). Approximately 5000 rpm are required to obtain a satisfactory particle size in a 500 ml flask. Sodium melts at 97.5° and heptane boils at 98-99° so that sufficient heat must be applied to the flask to keep the dispersion from coagulating. Once the dispersion was prepared, about 90 ml of tri—n-butylaluminum was added with a dropping funnel equipped with a pressure equaliz- Wp: .u-‘¢. Eli. “w s _ in. ing arm. The addition of the tri-n-butylaluminum took about an hour, after which the mixture was refluxed for two hours. The system was allowed to cool and then filtered through a fine frit filter stick into a clean round bottom flask. The filtrate was pale yellow and con- tained NaAl(But)u in n-heptane. The particle size of the aluminum metal that is formed during the reaction can be so small as to pass through a fine frit. If this happens the mixture should be allowed to stand for about a day, after which time the supernatant solution can be decanted. After the solids had been separated, the heptane was removed by vacuum evaporation. The final product was recrystallized from n-pentane at dry ice-acetone temperature. Care must be exercised here since the re- crystallization is not efficient and much product can be lost. Day reports that one recrystallization is 2A sufficient. However, the product prepared in this labora- tory was recrystallized twice. The melting point of 6A— 66° agreed exactly with that reported by Day (A7). The melting point sample was prepared by introducing a por- tion of the salt into a capillary tube which had been filled with dry nitrogen. A Hoover-Thomas melting point apparatus was used. The salt should be used shortly after it is pre— pared. It is difficult to prevent slow contamination with moisture even when the salt is stored under an inert atmosphere such as exists under normal dry box conditions. Sodium tetraphenylborate was obtained from the J. T. Baker Co., Phillipsburg, N.J. Rubidium and potassium tetraphenylborates were prepared from the sodium salt by adding a solution of the respective chloride to a solution of sodium tetraphenylborate. The resulting precipitates were washed with water until a negative test for chloride was obtained. Although it is easier to test for chloride than nitrate, it might be better to use the nitrate instead of the chloride since some diffi- culty was experienced in freeing the precipitate of all chloride ion. Nitrates usually are not adsorbed as much as chlorides. Sodium, ammonium and potassium tetrafluoborates were prepared by titrating a solution of the respective 25 carbonate with tetrafluoboric acid. The salts were crystallized and dried for use. The more common alkali metal salts such as perch- lorates, iodides, bromides, chlorides, etc., were ob- tained from commercial sources in the purest forms avail— able. Commercially available anhydrous lithium perchlorate has been found to contain as much as 8% water. The salt can be rendered anhydrous by heating to about 215° for about 2“ hours. The water content was checked by Karl Fischer titration. Lithium6 metal was obtained from Union Carbide's Oak Ridge Laboratory, Oak Ridge, Tennessee. The assay which was furnished with the metal showed 95.6% L16 and u.uz L17. In order to prepare L16 salts the metal was first added to water. Portions of the resulting basic solution were then neutralized with a reagent grade acid with the desired anion. The resulting salt solutions were crystallized, and the crystals were dried at 200°. Tetradeuteroammonium bromide was obtained from the Isomet Corporation, Palisades Park, New Jersey. The material was 98% isotopically pure and was used without purification. The lithium halides decompose tri-n-butylphosphate in the order; I>Br>Cl. The halides were simply dissolved in tri-n-butylphosphate. The resulting lithium di-n- butylphosphates_were filtered from the excess solvent and 26 washed thoroughly with acetone. The salts were then dried briefly at 100°. Preparation of Alkali Metal-1M2PY Solid Solvates The solid complexes Li(lM2PY)ClOu, and Li(lM2PY)SCN were most conveniently prepared by the following pro- cedure. A nearly saturated solution of the respective salt was prepared in nitromethane. Pure lM2PY was then very slowly added to the solution with stirring. The solvates readily come out of solution when the correct stoichiometry has been reached. The solid solvates were washed with cold nitromethane and then dried in a vacuum oven. The complexes Li(lM2PY)I, Li(lM2PY)Br, Li(lM2PY)Cl and Li(lM2PY)N03 were prepared by first dissolving the salt in lM2PY and then adding sodium dried benzene to the solution until the complex precipitated. The complexes were filtered and washed with cold benzene. All the lithium complexes are hygroscopic, and all manipulations should be carried out in a controlled atmosphere. The lithum-1M2PY complexes are white fluffy products. A suitable solvent for recrystallization of the complexes could not be found. Several polar and non-polar solvents were tried but even hydrocarbon sol- vents decomposed the complexes. The sodium iodide and thiocyanate solvates were prepared by cooling a warm concentrated solution of the 27 respective salt in lM2PY. The resulting solvates were filtered under vacuum, and the excess solvent was then removed by storing the solvates in a vacuum oven for about twenty four hours. Ammonium complexes NHu(lM2PY)NO and NHu(lM2PY)2SCN 3 were prepared in a manner similar to that of the sodium solvates. Efforts to prepare solvates with potassium, rubidium and cesium salts were not successful. Analyses All halogen and thiocyanate ions of the solid sol— vates were titrated with silver nitrate. Potassium chromate and acidic ferric ammonium sulfate were used as visual indicators (A9). Carbon, hydrogen and phosphorus analyses were per- formed by the Spang Microanalytical Laboratory, Ann Arbor, Michigan. All quantitative determinations of water were carried out by the Karl Fischer procedure (50). IV. RESULTS AND DISCUSSION Infrared Solvation Studies of Alkali and Ammonium Ions in 2—Pyrrolidones Preliminary investigations in the 4000 cm_1 to 80 cm‘1 infrared region indicated that 2-pyrrolidones ' would be convenient for solvation studies. The region E of most interest is from about 670 cm_1 to 80 cm-l. In 1M2PY, bands were found at 65A, 615, 557, A70, 308, and 175 cm-1. Katon §t_al, (51) examined many amides in the 700 to 250 cm_1 region. Their work on 1M2PY corresponds to the above values within experimental error. They did not observe the 175 cm‘1 band. Unfortunately, vibrational assignments have not been made for these bands. For the purposes of alkali ion solvation, however, we are more interested in the availability of spectral windows. Since lM2PY exhibits reasonable transparency below A50 cm—l, it was suspected that the solvent would be useful in studying solvation phenomena. As stated earlier, 1M2PY has a dielectric constant of 32 and a dipole moment of H.09 D and is a good dissociating solvent. There is a striking difference between the pure l-methy1-2-pyrrolidone spectrum and the spectrum of 1M2PY which contains a lithium salt. Figure 1 illustrates 28 Figure 1.--Far infrared spectra. A. Pure l-methyl- 2—pyrrolidone. B. 0.3 fl LiClOu in l-methyl- 2-pyrrolidone. iiiiiii 31 the difference. Spectrum A is of the pure solvent, and spectrum B is of solvent which contains 0.3 M anhydrous lithium perchlorate. As illustrated in the figure there is a new band present in the lithium perchlorate solu- tion. The spectra shown in the figure were obtained with air in the reference beam. The exact position of the band maximum can be ob- tained by subtracting out the absorption which is due to the solvent. The tedious work of baseline subtractions can be avoided in this case if an equal thickness of solvent is placed in the reference beam. The use of solvent in the reference beam produces a very flat base- line. In any case, the new band maximum was measured at 398 cm.1 with an uncertainty of i Hem-l. The new band is not due to the perchlorate anion. Perchlorate has a band at about 620 cm.1 (52), but it does not interfere with the 398 cm"1 band. Furthermore, for lithium, the band occurs at exactly the same frequency regardless of anion, except for the chloride. The entire spectral region from the ultraviolet to the practical working limit in the far infrared was examined. No other bands characteristic of the solutions could be found. Similar experiments were carried out using 2-pyrrolidone and l-vinyl-2-pyrrolidone. As stated earlier these solvents have physical prOperties similar to l-methy1—2-pyrrolidone (see Table A). The results are tabulated in Table 5. As seen in the table, 32 the band frequencies change slightly from solvent to solvent but the anion independence is very obvious in all three solvents. TABLE 5.-—Absorption spectra of lithium ions in 2- pyrrolidones. v max. cm"l Salt lM2PY 2PY lV2PY LiI 398 t A uoo : u 419 i u LiBr 398 uoo A19 LiCl 377(398)* uoo M19 LiClOu 398 Ace 419 LiNO 398 397 “19 L180 398 uo2 A19 *Small shoulder at 398 cm-1. Other alkali metal and ammonium salts were examined 1 to 80 cm'l. in lM2PY, 2PY, and lV2PY from A000 cm- Sodium, potassium, rubidium, and ammonium ions also pro- duce anion independent infrared bands. The results are tabulated in Table 6. Solvent opacity prevents the study of lV2PY solutions in regions where the heavier alkali ions absorb in the other solvents and thus only data for lithium salts could be otained. It is obvious from Table 6 that the band maximum for a given cation depends on the cation and the solvent but not on the anion. The anion independent nature of the bands strongly suggests a 33 TABLE 6.--Absorption spectra of alkali metal ions and ammonium ions in 2-pyrrolidones. v max.(cm'l) Salt 2PY 1M2PY 1V2PY LiI A00 i A 398 1 A A19 i A LiBr A00 398 A19 LiCl A00 377(398) A19 LiClOu A00 398 A19 LiNO3 397 398 A19 LiSCN A02 398 A19 NHuI 216': A 206 i A NHuBr 217 207 NHuClOu 218 207 NHuNO3 218 213 NHHBFA 218 207 NHuHSOu —-— 207 NHuB(¢)u 218 --- NHuSCN 219 207 NaI 207 i A 20A i A NaBr 206 20A NaClOu 207 20A NaNO 207 206 NaSC 206 207 NaB(¢)u --- 205 NaBFu 206 207 Kl 1A7 i 5 1A0 i 5 KClOu 1A5 1A0 KSCN 1AA 138 KB(¢)u 1A7 138 RbI --- 106 i 6 RbClOu —-— 106 RbB(@)u --- 106 3A cation-solvent interaction. The bands, as they appear in 1M2PY, are shown in Figure 2. It is seen from the data in Table 6 that the fre— quencies of the bands are dependent only on the nature of the cation and decrease progressively with the increase in the atomic number of the metal ion. In the case of the ammonium ion, however, the frequency of the cation band in 1M2PY is about equivalent to the frequency of the sodium band, while in 2PY, the frequency is higher than expected from the mass effect. The absorptivities of the different cations pro— gressively decrease with increasing atomic number of the cations. Here again the ammonium band does not fall into the general pattern, being approximately equal to the intensity of the lithium band (Figure 2). The positions of the far infrared bands do not change upon the addition of small amounts of water (~l%). The only noticeable effect is the upward shift of the baseline due to water absorption. Since most of the measurements were done with 0.5 M_solution of salts and the water content of the solvent was less than 0.011% it is seen that even if water had greater cation solvating power than the pyrrolidones, which is not likely, the effect of water would be minimal. Several questions about the nature and origin of the bands immediately arise. Is there a definite 35 Figure 2.--Far infrared absorption bands of 0.5 M alkali metal ions in l—methyl- 2-pyrrolidone. 36 JONVEUOS‘GV 37 stoichiometric solvent structure around the solvated cation? Does the cation interact with the solvent mole— cule at a definite sight in the solvent molecule? If the alkali ions interact strongly enough with the solvent to produce discreet infrared bands, is the interaction strong enough to allow the isolation of salt-solvent adducts? What is the role of the anion? What is the nature of the bonding between the alkali ions and the solvent? Do the cation-solvent bands obey Beer‘s law? These and other questions must be answered before an intelligible picture of solvated ions can be put to— gether. The following different combinations of ions and solvent can be visualized: l. Ions completely surrounded by solvent with no or very little interionic interaction. 2. Ions completely surrounded by solvent but interacting to form ion pairs. 3. Ions paired but sharing solvent. A. Intimate or contact ion pairs. 5. Formation of a pseudo lattice structure in solution. It is known (53, 5A) that the band frequencies of lattice modes of vibration are anion dependent. For example, in going from LiF to LiBr the frequency of the transverse lattice vibration changes from 307 cm“1 to 38 159 cm’l. If the solution bands shown in Table 6 were of the same nature as lattice modes, a more pronounced anion dependence would be expected. Likewise, if ion pairs of one form or another were responsible for the solution bands, their frequencies would probably be affected by a change in the anion. It is also conceivable that a hindered solvent rota- tion, not obvious in the liquid phase, might become more noticeable in the presence of a solute. If this effect were the cause of the bands which appear in alkali salt solutions, it would be expected that the band frequency would be independent of cation, or nearly so. It is apparent that the alkali ions give rise to infrared bands of widely different frequencies which depend on the nature of the cation. The data illustrated in Table 6 are most easily explained if it is assumed that solvent molecules completely surround the cation and that the observed bands are due to ion—solvent vibrations. It is worthwhile at this point to comment on the meaning of the word complex as it will be used in the context of this discussion. Experience has shown that semantics problems can cause undue confusion. This thesis deals almost exclusively with solutions of alkali and ammonium ions.~ Solvated species which contain these ions will be considered to be complexes. The word 39 complex is not meant to automatically denote covalency, and conclusions about the nature of the bonding should not be assumed. The absorbance of the LiClOu—1M2PY band was studied as a function of concentration. The results are shown in Figure 3. Small solvent absorption corrections were made for each sample. The solution spectra of samples varying in concentrations from about 0.1 to 1.0 molar were obtained in order to determine if the band maximum was a function of concentration. The band maximum 1 for all solutions. There are two occurred at 398 cm- major problems in obtaining quantitative absorption data for the lithium band. First, integrated absorbances are difficult to obtain because of the close proximity of solvent bands. Second, if the whole band is scanned and absorbance is measured at a given wavelength, instrument instability can introduce considerable error if a large time lag exists between the first and last sample spec- trum. In order to minimize the time required to obtain absorbance readings, the following method was used. The monochromator was set at a fixed wavelength, such as the maximum of the analytical band. A sample was placed in the beam and the recorder pen was allowed to come to equilibrium. The reading was recorded. The cell was then flushed with acetone and pentane and dried without dismantling as described in the Experimental Part. AO Figure 3.--Absorbance vs. concentration for LiClOu in lM2PY. Path length = 0.0 mm. Absorbance readings taken at 398 cm- . Al Aflv COHpmppsoocoo To me to ma me 11.0 me «.0 _ _ a a _ _ _ _ m mhzmflm m.o m.o eoueqaosqv A2 Several readings at different concentrations can be ob- tained in this way in only a few minutes. It is seen in Figure 3 that the LiClOu-1M2PY band follows Beer's law to a remarkable degree. This verifies the fact that the 1 absorption at 398 cm" definitely is caused by the presence of LiClOu. The data that accompany Figure 3, as well as absorption-concentration data for NaClOu in 1M2PY, are tabulated in Table 7. TABLE 7.--Beer's law data for LiClOu and NaClOu in 1M2PY. Solution No. Concn. fl, Absorbance at 398 cm"1 a LiClOu 1. 0.97 0.5A 2. 0.7A O.Al 3. 0.56 0.33 A. 0.38 0.22 5. 0.2A 0.12 5. 0.085 0.05 b NaClOu l. 0.62 0.35 2. 0.51 0.26 3. 0.Al 0.23 A. 0.3A 0.19 5. 0.17 0.11 aPath length of 0.05 mm. bPath length of 0.1 mm. A3 If the cation is truly involved in the pr0posed alkali ion-solvent vibration,then a cation isotopic substitution should produce a shift in the band fre- quency. Lithium7 isotope produces band maxima at 398 cm'1 independent of anion except for chloride. The 1 salts which produce maxima at 398 cm- include the iodide, bromide, nitrate, perchlorate, thiocyanate, and tetra- hexylaluminate. The lithium6 salts studied are tabulated in Table 8 along with the band maximum for each salt. In going from lithium7 to lithium6 the mass change is about 1 for lithium7 salts, a shift of 22 cm'l. 13.5%. The change in energy is from 398 cm- to A20 cm"1 for lithium6 TABLE 8.--Absorption spectra of Li6 salts in 1M2PY. Salt Ymax LiClOu A20 1 A LiNO3 A20 LiI A20 LiBr A20 LiCl A09 1 6 A semi—quantitative calculation can be made to predict the magnitude of the expected shift. Sample calculations are included in Appendix II. Since only AA one cation-solvent vibration has been observed for each cation, the simplest model would involve the cation vibrating against one molecule of solvent. In this case, the calculated value is A26 cm-1 1 , differing by only 6 cm" from the experimental value. The usual assumption was made that the force constant does not change in the cation isotopic substitution. Experimental results, which will be discussed later, indicate that at least four coordinated molecules of solvent are required in order to observe the 398 cm.1 band in 1M2PY solutions of lithium salts. A calculation of the expected isotope shift for a four coordinate tetrahedral model can also be made (see Appendix II). In this case the assumption must be made that the v” low energy deformation mode does not change much in going 6 1 from Li7 to L1 and a value of A27 cm' was calculated for the frequency of the Li6-1M2PY band. This value is only 1 cm-1 different from the diatomic case. Regardless of which model is used, the diatomic, the linear triatomic or the tetrahedral, the calculated values lie within a few wavenumbers of one another (see Appendix II). A similar type of isotopic substitution experiment was carried out with NHIQ and ND+u ions. In this case the bromide salt was used. If it is again assumed that the cation vibrates against one solvent molecule, the calculated value of the band frequency is 191 chI. The A5 1 1 experimental value is 192 cm’ i 3 cm' , a remarkably good agreement. The most important result of the isotOpe experi- ments is that they substantiate the fact that the cation is involved in the far infrared vibrations. This idea is, of course, in close agreement with the anion inde- pendent nature of the bands. 6 and L17) produces a Lithium chloride (both Li band at lower frequency than any of the other lithium salts. Unfortunately, the heavier alkali ion chlorides are not soluble in 1M2PY so that comparisons cannot be made. The exact reason for this discrepancy cannot be offered at this time. It may be more than coincidence that the chloride is the smallest and lightest of all the anions which were studied. It is conceivable that the relatively small mass and size of the chloride ions leads to the formation of ion pairs and that the fre- quency of the cation vibration is thus affected. This explanation is not entirely satisfactory since the lithium chloride result does not deviate from those of the other lithium salts in 2PY or 1V2PY. Furthermore, both of these solvents have lower dielectric constants than lM2PY. It has been suggested that anions are more highly solvated in protic solvents than aprotic solvents (55). Since 2PY is protic and lM2PY is aprotic, different A6 types of anion solvation could produce the difference in chloride salt results between the solvents. After it was once established that cation-solvent vibrations were responsible for the far infrared bands, the question of stoichiometry around the cation seemed most important to answer. There are many spectrosc0pic techniques available to determine the coordination number of complexes (56). The application of any of the methods to the infrared study of alkali ion complexes in solution meets with at least two difficulties. First, quantitative experiments are difficult to perform in the far infrared region where the alkali bands occur. This difficulty arises from such factors as low molar absorp- tivities, band overlaps and baseline drift. Second, a solvent must be found that does not successfully compete with the complex ligand and which allows the ratio of ligand/metal to be varied over a wide range. This range must necessarily include ratios between one and ten, since most coordination numbers fall in this range. In the case of polar solvents such as the pyr- rolidones, it seems that ionic solvation results from an ion-dipole interaction between the solvent molecules and the ions in question. In order to study pyrrolidones- alkali metal ion interactions,an "inert" non-polar solvent which would not compete with the pyrrolidones would be ideal in which to carry out a mole ratio experiment. In A? an attempt to find such an inert solvent in which the 1M2PY/LiClOu ratio could be varied appropriately, many solvents were investigated. These included benzene, cyclohexane, heptane, nitromethane, alcohols, halo- hydrocarbons, dioxane and acetonitrile. For the study of the Li-lM2PY band, either low solubility or undesir- able spectral prOperties rule out all these solvents with the exception of dioxane. Lithium perchlorate by itself is only slightly soluble in dioxane but with small amounts of 1M2PY present, a 1M2PY/LiClOu ratio of 1.25/1 can be attained at a salt concentration of 0.5 M. A range of solutions with ratios varying from 1.25 to 10 were prepared. The spectral region from A50 cm'"1 to 325 cm-1 was examined for changes in the usual 398 cm-1 band as it appears in pure 1M2PY. All spectra for the infrared mole ratio study were obtained with a blank of appropriate lM2PY-dioxane composition in the reference beam. In going from 10 R to A R (R = lM2PY/LiclOu ratio) 1 the 398 cm' hand first increased slightly in frequency 1 to A08 cm'l. Beyond A R the A08 cm" band decreased in intensity, broadened and the apparent maximum shifted until the limiting value was reached at 2 R and 368 cm-1. An attempt was made to measure and correlate variations in intensities of either band with the LiClOu mole ratio. This was not possible because of the close proximity of l 1 the 398 cm' and 368 cm' bands. However, the shift in A8 1 to 368 cm"1 could be the band maximum from 398 cm- followed with a reasonable degree of accuracy. A plot of the band maximum vs. lM2PY/LiClOu mole ratio produced the results shown in Figure A. More important than the slight increase in frequency between 10 and A R is the drastic change in the nature of the band at A R. There is a definite break in the plot at AR. This is inter- preted to mean that four moles of 1M2PY per mole of lithium ion are required to maintain the structure which 1 band in pure 1M2PY. is responsible for the 398 cm- A coordination number of four is very common for lithium ion in solids (57, 58) and, therefore, a solva- tion number of four is not too surprising. Scale molecu- lar models show that four molecules of 1M2PY can easily fit around a lithium ion. Although the solid Li(lM2PY)C10u complex prepared in this laboratory contains only one molecule of 1M2PY, it is interesting to note that Li(lM2PY)uClOu'H2O which contains four molecules of 1M2PY has been prepared by another investigator (2A). As R falls below four, either a new 1M2PY complex is formed such as L1(lM2PY)x+ or the 1M2PY molecules in the solvation shell are replaced by dioxane to give species such as Li(lM2PY)x(Diox)y+. If the physical properties of dioxane are considered, it becomes obvious that a large degree of ion pairing is also possible at low R values. Pure dioxane has a dielectric constant of A9 Figure A.--Frequency of the solvent—cation absorption band vs. the lM2PY/LiClOu mole ratio in dioxane solutions. LiClOu concentration = 0.5 M. : mhswflm wo_o.+_._ >QN2. 0.5.41 1 A» 50 _ _ _ _ _ _ 0mm J own poem I 00m 100m looe lomv Lone WflW/XVW JO HBHWHNH/WM 51 2.2 (59), and a dipole moment of 0.A5 D (60). In 0.5 fl LiClOu solution the anion could very well be involved in 1 band. the solvate that is responsible for the 368 cm- As mentioned earlier, lithium perchlorate is only very slightly soluble in pure dioxane. The limited solubility,along with the large increase in solubility when small amounts of 1M2PY are added, suggest that a pure Li—dioxane solvate does not exist in the presence 1 band were due to of 1M2PY. Conversely, if the 368 cm— a pure Li-(1M2PY)X+ solvate of stoichiometry less than A:l but greater than 1:1, the absorbance of the band should decrease rapidly as R approaches one. This does not happen. Also, it would be expected that a pure lower order Li-(1M2PY)+ complex would absorb at higher energy instead of lower energy. This would be expected because as R decreases the lithium ion is shared by a lesser number of 1M2PY molecules and the Li-1M2PY bond would most likely become stronger. In View of the above evidence, it is apparent that the 368 cm_1 band is probably due to a composite species made up of lithium ions, 1M2PY, dioxane and perchlorate ions. Unfortunately, this study could not be extended to other alkali ions because of their limited solubility. 52 Correlation of Ion-Solvent Band Frequencies with Physical Properties of the Ions For a diatomic system, Badger's rule relates the force constant to the internuclear distance according to the following equation (61, 62): k = 1.86 x 105 There k is the force constant in dynes/cm, r is the internuclear distance in Angstroms and diJ is a constant characteristic of all molecules made up of one element in the ith_row and one in the Jth_row of the periodic table. Similar relationships have been sug- gested by Clark (63) and also by Allen and Longair (6A). In applying such considerations to the far infrared ion-solvent bands it was found that a radius dependence invariably excludes ammonium ion from any trend among the alkali ions. The crystallographic radius of ammonium ion falls between that of potassium and rubidium (65), whereas the ion-solvent vibrational frequency is slightly above that of sodium (see Table 6). Hooke's law is a good approximation for many diatomic systems and is stated as follows (66): --_1_ .11 v ' 2wc u 53 There 5 is the frequency of the vibration in cm"l, f is the force constant in dynes/cm, c is the Speed of light m182 in cm/sec. and u = 51:53 where m1 and m2 are the re- spective atomic masses in grams. As stated earlier (see Figure 2), the bands of the alkali ion vibrations exhibit a regular decrease in the frequency and intensity as the mass of the cation in- creases. If a diatomic model consisting of a cation and solvent molecule is considered, it becomes obvious that the mass of the cation and the force constant for the ion—solvent bond will determine the frequency of the vibration. The force constants of the ion-solvent bonds are not known, but it may be possible to relate them empirically to some specific physical property of the cation. It appeared reasonable to assume that the force constant for the ion-solvent vibrations in a given solvent may be related to the reciprocal of the cation mass. Under this condition, assuming a diatomic model, the frequencies of the vibrational bands in a given solvent will be given by the following formula. ¢rkproportion ///mcation + Insolvent V - 2nc m cation msolvent 5A The terms have their usual significance. With isotopes, the force constant is assumed not to change and the cal- culation is based on a change in reduced mass only. If this relationship has any real meaning, a trend should in +m be observed by plotting 5 vs. a——l——— /// cation solvent cation solvent The results of such a plot for the alkali and ammonium ions in l—methyl-2-pyrrolidone and dimethylsulfoxide are shown in Figure 5. It would probably be somewhat naive to expect that calculations of this type would be successful for all ion-solvent interactions. Nevertheless, it is obvious from Figure 5 that such considerations may be useful for predicting frequencies that can be expected from a given cation. The relationship between force constant and mass suggested above is similar to the relationship between force constant and radius as suggested by Badger. It should be noted that both relationships are empirical. Infrared Anion Studies Several papers have appeared recently which discuss the vibrational spectrum of nitrate ion in ionic media (67-69). The nitrate ion spectrum serves as an indi- cator of specific ionic interactions. These interactions become evident from the alteration of the nitrate ion symmetry with a resulting change in its vibrational spectrum. 55 L m + m Figure 5. ___;;___ cation solvent Ki cation solvent wavenumber of infrared band maxima for alkali and ammonium ions in 1M2PY and DMSO. 56 m mhsmfim H 5 a 03 co: com com _ l , 1 omza U .1 is? O .1 m2 IL 152 IL emz J J J .IIL IIJ l a III was .I OH om om on om om on om om ooa OHH oma ems . 03H oma 00H oma uotieo m1 ueA os A I m quehtosm + uotieo EDI X 57 An attempt was made to check for a possible anion dependence on the alkali ion-solvent bands by the addi- tion of excess anion in the form of tetralkylammoniom salt. Tetralkylammonium ions are considered not to be highly solvated and therefore would not be expected to interfere with the alkali ion solvation. Nitrate, chloride and bromide anions were added to lithium salt solutions in 1M2PY. The results are listed in Table 9. No frequency or intensity changes outside experimental error were observed in either the cation-solvent band or the nitrate fundamentals. If it can be assumed that cation-anion interaction would distort the cation-solvent interaction or the nitrate symmetry to produce changes in their respective ir fundamentals, then no significant cation-anion inter— action is indicated. The conductance data of Dyke et_al. (8) show that the ion pair formation constants of sodium perchlorate, iodide and tetraphenylborate are all zero, since the data fit the Fuoss-Onsager equation (33, 3A) for un- associated electrolytes. The salt concentrations for the study were on the order of 10"3 molar. Study of the Carbonyl Frequency in 1M2PY—Dioxane Mixtures In addition to a change in the Li—lMZPY vibration with changing lM2PY/Li ratio, a considerable difference TABLE 9.--Infrared absorption data: 58 anion study. Salt Concn. g v2 N03- V3 NO ' vL1-1M2PY LiClOu 0.1 -—- --- 399 cm"1 ButuNNO3 0.6 835 13A3 ——— ButuNNOB 0.05 too 13A5 ——— weak ButuNNO3 .6 83A 13A3 397 1101024 0.05 ButuNNO3 .6 837 13A3 399 LlcloLl .6 But NNO .05 too l3A7 A 3 weak LlcloLl .6 399 LiClOu 0.2 398 cm'1 LiClOu 0. 398 (But)uNCl 0 LiCl 0.2 377 Lil 0.1 398 cm"1 LiI 398 cm—1 ButuNI 0.8 59 is also noticed in the frequency of the carbonyl band as illustrated in Figure 6. In this study the lM2PY con- centration was held constant at 2 M. In 2 molar 1M2PY in dioxane the carbonyl stretch occurs at 1692 cm-1. At low lM2PY/Lil values (~2) the v 1 C=O band has shifted At A R the band is split with maxima at l to 1669 cm- 1692 and 1669 cm‘ When R is 10, the maximum falls at 1692 cm_1 and a definite shoulder is apparent at 1669 cm-1. The 1692 cm"1 band is assigned to the un- complexed carbonyl band. The appearance of the new lower energy branch when lithium iodide is present is assigned to the 1M2PY complexed with the lithium salt. Replacement of lithium iodide with lithium perch- lorate produces identical results. The results are also the same regardless of whether the ratio is varied by changing the ligand or the salt concentration. Since a similar shift can be observed upon the addition of water, a rigorously anhydrous system was used in the study. The carbonyl band is also split in pure 1M2PY which is prob- ably a result of inter—solvent interaction. It should be noted that in examining the VC=O band of 1M2PY as a func- tion of 1M2PY concentration in dioxane it eventually splits as the concentration of 1M2PY increases and approaches pure 1M2PY. This type of interaction is not a factor inihe results depicted in Figure 6 because the 60 Figure 6.--The carbonyl band of 1M2PY at various ratios of lM2PY/Lil. Concentration of lM2PY = 2 M. Lithium iodide concentration variable. 61 com. O. «K 000. \ ____l_ cum Nam .120 com. com. com. com. _J__fi.fij jun—ea. g m mhsmflm w 2365.12. >mN2_ 2N 00m. 000. [Ill 1:: _d—__ ‘ 62 1M2PY concentration is constant and low enough to pro- duce only a single band. Examination of the 1M2PY spectrum in the A000- 600 cm"1 region showed that, with the exception of the carbonyl band, the spectrum is not noticeably affected by the addition of lithium salts. In view of the fact that the v 0:0 splits signifi cantly in the presence of lithium salts, it is reason— able to assume that the ion—solvent bonding occurs at the oxygen atom. In both transition metal (26) and non- transition metal (2A) solid complexes with pyrrolidones, the oxygen atom rather than the nitrogen has been shown to be involved in the bonding. It has been postulated that amides can be represented by the following three valence- bond structures (70). II III The lowering of the carbonyl stretch indicates that the carbon-oxygen bond becomes weaker upon complexation. Lowering of the carbonyl bond order such as in I or III would be consistent with the lowering of the carbonyl 63 stretch as is observed in complexation. Therefore it is possible that complexation favors structures I or III or both. Solid Solvates of Alkali Salts with lM2PY During the search for a suitable solvent in which mole ratio studies could be carried out, the nitromethane lM2PY—LiClOu system was investigated. Nitromethane does have a relatively high dipole moment, 3.17 D (71), but it is a poorer donor than lM2PY and was considered to be a good prospect. The solubility of lithium perchlorate in nitromethane is considerable. In order to test the miscibility of the three component system, a small amount of lM2PY was added to a concentrated solution of lithium perchlorate in nitromethane. Long, needle-like and very well defined crystals began to grow in the solution. The crystals were filtered, washed with nitromethane and dried in a vacuum oven at about 35°. The carbon-hydrogen analysis indicated a Li(lM2PY)ClOu complex. Similar complexes with lithium thiocyanate and lithium iodide can be prepared in the same way. Un- fortunately, the chloride and the bromide are not soluble in nitromethane. The bromide and chloride complexes can be prepared by adding benzene to a concentrated solution of the salt in lM2PY. Sodium and ammonium solid solvates were also prepared by cooling warm concentrated solutions 6A of the salts in lM2PY. The results are tabulated in Table 10. Attempts to prepare solid solvates with potassium, rubidium and cesium salts were not success- ful. Since 1M2PY contains both oxygen and nitrogen there is a possibility of coordination taking place at either sight. The infrared spectra of all the complexes were obtained. It was discovered that in all cases the carbonyl band shifted to lower frequencies. The fre- quencies of the vC=O band for the solids are shown in Table 10. It is seen that for the lithium salts, the 1 band shifts to lower frequency by about A2 cm' and that the shift is completely independent of anion. It should be noted that the band shift is considerably greater than the one observed in solutions showing that in the solid complex there is a stronger interaction between the carbonyl group of 1M2PY and the lithium ion than in solution. On the other hand, in the case of sodium and ammonium complexes, the magnitude of the band shift seems to be dependent on the nature of the anion. It should be observed that in the case of the sodium and ammonium solvates the anion is not the only variable. In addition to the difference in the anion there is also a difference in the number of coordinating lM2PY mole— cules. It is likely that the difference between the 65 TABLE lO.—-Solid complexes of 1M2PY with alkali and ammonium salts. Analyses Salt Calcd. Found vC=O 0 H MX 0 H MX 1M2PY -- -- —- -- —- -- 1687 Li(lM2PY)ClOu 29.22 A.A1 -- 29.31 A.3o -- 16A5 Li(lM2PY)SCN -- -- 39.61 —— -- 39.08 16A5 Li(lM2PY)Cl —— -- 29.95 -- -- 29.A8 16A5 Li(1M2PY)Br 32.29 A.89 —- 33.15 5.79 -— 16A5 Li(lM2PY)I -- -- 57.A5 -- -- 57.52 16A5 Na(lM2PY)3I -- —- 33.51 -— -- 33.25 1655 Na(lM2PY)2ClOu 37.A5 5.65 —— 38.38 6.05 -- 1670 NHu(lM2PY)NO3 33.52 7.31 -- 3A.62 7.70 —- 1658 NHu(lM2PY)2SCN -— -- 27.7A -- -- 26.66 16AA 66 vC=O for the sodium complexes, or ammonium complexes, is due to the difference in stoichiometry as well as to the difference in the anion. The far infrared spectra of the solids were usually not well defined in the regions where the cation-1M2PY bands occur. Reproducible sharp bands at 375 cm'"1 and 395 cm"1 were observed for Li(lM2PY)Cl and Li(lM2PY)SCN complexes respectively. In the case of the ammonium and sodium complexes, only very broad bands were observed in the general region of the corresponding cation-solvent bands obtained in solution. In the 200 cm‘1 region lM2PY itself exhibits a gradual increase in absorbance which complicates the solids spectra. Raman and Ultgaviolet Spectra of Alkali Ion Solutions Attempts were made to study the solvation of alkali metal ions in 1M2PY by ultraviolet and Raman spectro- scopies. The ultraviolet spectrum of 1M2PY is feature- less from 380 mu to the cutoff at 265 mu. Spectra of solutions of 1.0 M_lithium perchlorate, 0.6 M_sodium tetraphenylborate and 0.8 M_sodium perchlorate were obtained from 380 mu to 265 mu with negative results. The sodium tetraphenylborate solution exhibited absorb- tion due to the anion. 1 Raman spectra in the 1700 cm- to 200 cm"1 region were obtained of 1.0 M_solutions of lithium chloride, 67 lithium bromide and lithium perchlorate in lM2PY. Three anions were chosen in order to observe possible anion effects. Particular attention was paid to the A00 cm-l region, since the infrared lithium band occurs at 398 cm"1 in 1M2PY. The perchlorate v2 deformation mode was ex— hibited at A55 cm-l. This band was not present in the chloride or bromide solutions and was due to the per— chlorate anion. No bands which could be attributed to salt-solvent interaction were discovered. The absence of Raman bands seems to indicate that the ion-solvent interaction is largely ionic. The alkali salts are ionic and 1M2PY has a high dipole moment (A.09.D% The nature of the salt-solvent systems studied and the above Raman data suggest that the ion-solvent bond is formed by the ion-dipole interaction. The idea of relatively weak ion—dipole forces is compati- ble with low infrared frequencies, as are observed with the alkali ion vibrations. It should be noted that covalency among the alkali metals is not unknown. Li2 exists in vapors of the metal and the Li—Li bond has been shown to be about lAz p character (72). Nuclear Magnetic Resonance Study of Alkali Ion Solvation The chemical shift of a proton is very dependent on its environment. If the electron distribution in the 68 proximity of the proton is perturbed, the shielding of the proton is changed and the chemical shift of the NMR resonance is altered. With this thought in mind, it is interesting to consider the electronic configuration of the amide moiety in lM2PY. The 2p w orbitals of the OCN group are de- picted as follows (73). \ The oxygen, carbon and nitrogen 2p w orbitals con- tain four electons. Maximum overlap exists, especially in the planar 1M2PY, to form a n molecular orbital which embraces the OCN group. It is reasonable to assume that if lM2PY were to) act as a ligand, or as a strong solvating agent, the electronic shielding characteristics of the system would be perturbed enough to produce a change in the chemical shift of the protons which are near the n system. Three mole-ratio studies on the pyrrolidons-Li+ system have been carried out using proton nmr spectra. In the first study, the position of the N-methyl protons 69 resonance of 1M2PY was followed as a function of the lM2PY/LiClOu ratio. The concentration of LiClOu was varied so as to produce mole ratios from 1.0 to 9.5. The concentration of 1M2PY was held constant at 2.0 M, Figure 7 shows the results. There is a definite break at a ligand to salt ratio of A.3. In the second nmr eXperiment, the LiClOu concen— tration was held constant and that of 1M2PY was varied. The methyl protons resonance was again plotted as a function of the ligand to salt ratio. Figure 8 shows that the value of A.5 for the coordination agrees quite well with the value of A.3 obtained by varying the salt concentration. The baseline illustrates that the posi- tion of the methyl protons resonance is not a function of the 1M2PY concentration. A third mole ratio experiment was performed using lithium iodide instead of lithium perchlorate. The difference in the nature of the anions could produce a change in the coordination number value if the cation is involved in the coordination sphere of the lithium ion in solution. The salt concentration was varied and the results are shown in Figure 9. Although more curvature is evident in this case,the coordination number value of A.5 agrees well with the values obtained previously. Table 11 summarizes the results which include the far infrared mole ratio results which were discussed 70 Figure 7.—-Chemical shift of the N-methyl protons of 1M2PY gs: lM2PY/LiClOu mole ratio. Concen— tration of LiClOu varied. Concentration of 1M2PY = 2 M. 71 38...: .1. Ass >n_Ns__ O_._.aNs: 0.94m o_ m m h m n v m N _ .7 _ a . _ _ _ _ a . nausea . 0: :L N: mt .1:— m: 0.: NC mt SdO 76 earlier. It is seen that all the data indicate a solva- tion number of four for the lithium ion in 1M2PY solu- tion. It is possible that the deviations from the integral value of A are due to solvation of the anion. TABLE ll.--Coordination number of Li ion in 1M2PY. Method Salt Variable Coord. No. nmr LiClOu LiClOu A.3 nmr LiClOl4 lM2PY A.5 nmr LiI LiI A.5 far ir LiClOu 1M2PY A.2 The reasonance of the dioxane protons was also studied as a function of lithium perchlorate concentra— tion. The results are shown in Figure 10. At low ratios of lM2PY/LiClOu, the dioxane protons undergo an appreciable shift. However, there is no dioxane proton shift without salt present. It is, therefore, possible that at low concentrations of 1M2PY, dioxane participates in the solvation of the lithium ion. These results also support the hypothesis that the solvated species which gives rise to the 368 cm"1 band in the far infrared mole ratio experiment may include dioxane in its solvation sphere. The absence of a shift in either the dioxane 77 Figure 10.--Chemical shift of the dioxane protons y_s_. lM2PY/LiClOu mole ratio. 1M2PY concentra- tion = 2 M. OH ohswflm 168...: >mm2. 02km 78 m n v m _ _ a .ONN .NN NNN ‘ MNN VNN .VNN 5:10 79 protons or 1M2PY protons in binary solutions of the two solvents is consistent with the idea of minimum dioxane- 1M2PY interactions. The sodium-1M2PY and sodium-DMSO (l, 3) infrared bands, as studied in this laboratory, have previously been described in detail. Little substantial informa— tion exists on the number of solvent molecules associated with the sodium ion in 1M2PY or DMSO. Schaschel and Day (22), however, recently determined the number of tetrahydrofuran molecules associated with sodium ion by a proton magnetic resonance method based on chemical shift.‘ Their results indicate that THF and sodium ion form both a A:1 and 1:1 complex in solution. They used sodium tetrabutylaluminate as a source of sodium ion. For our purposes,the advantage of NaAl(But)u is that it is soluble (~0.5 M) in organic solvents such as heptane, cyclohexane and dioxane and, therefore, it was used in proton magnetic resonance mole ratio studies which are similar to those already described for lithium. Since NaAl(But)u is soluble in hydrocarbons, it would be desirable to use a hydrocarbon as an inert sol— vent for a mole ratio study. Unfortunately, 1M2PY, NaAl(But)u and hydrocarbons such as cyclohexane, benzene or heptane are not miscible in the pr0portions of interest. Similar comments apply to DMSO, NaAl(But)Ll and hydrocarbon systems. 80 As illustrated earlier, dioxane does not appre- ciably interact with 1M2PY. It also has been recognized by other workers that dioxane is a poorer solvating agent than DMSO or amides (7A, 75). Dioxane would not be ex- pected to hydrogen bond strongly or solvate anions as would a protic solvent (55). As in the study of lithium solvation in mixed solvents, dioxane was again chosen as an "inert" solvent in which to study sodium solvation with 1M2PY and DMSO. Figure 11 shows the mole ratio result for DMSO. The two straight parts of the curve can be extrapolated to meet at 6 R (R DMSO/NaAl(But)u). The shift of the DMSO protons above 6 R is largely determined by the excess DMSO which is present. The proton shift below 6 R is apparently determined by DMSO interacting with either NaAl(But)u or dioxane or both. The baseline shows the proton shift with no salt present. Dioxane and DMSO do show some interaction as illustrated in the baseline in Figure 11. Consideration must also be given to the fact that TMS itself is not completely unaffected by the change in DMSO concentration. The additive interactions of DMSO and dioxane with respect to TMS is reflected in the baseline. The same composite effect on the DMSO protons should be present in Curve A, Figure 11, along with the effect of the NaAl(But)u. In any case, Curve A, Figure 11, shows that the salt has the most influence. 81 Figure ll.--A. Chemical shift of DMSO protons vs. DMSO/NaAl(But)u mole ratio; NaAIIBut) concentration = 0.36 M. B. Blank. ICPS 82 I57 I56 7 l5!) I54 I53 F I52 P I5I — I I I I L I J I I I I I I I I I I I 0 I 2 3 4 5 6 7 8 9. IO II I2 I3 I4 I5 I6 I7 I8 ' RATIO DMSO No+ Ausm); Figure 11 83 Since the baseline does not have zero slope, a correction to the mole ratio plot should be made. In order to make an accurate correction, the dissociation constant of the complex must be known so that a correction can be made for the uncomplexed DMSO. If the whole baseline value were simply subtracted, a maximum correction would be affected. The result of subtracting the baseline is shown in Figure 12. The value of 6:1 at the break point does not change. Only the shape of the mole ratio plot changes. The break at 6:1 is taken to mean that after 6 molecules of DMSO are available for every mole of NaAl(But)u the cation influences the DMSO resonance to only a minor degree. If the assumption is made that the bulky A1(But)y-anion is solvated to a small extent com-7 pared to the cation, then the solvation number of sodium ion in DMSO is most likely 6. Figure 13 shows the results of the l-methyl- 2-pyrrolidone nmr mole ratio eXperiment. The break occurs at A.3:l mole ratio which implies a solvation number of A for sodium ion in 1M2PY. In this case the baseline does have zero slope and no corrections to the mole ratio line are necessary. The N-methyl protons resonance was used rather than that of the methylene protons adjacent to the carbonyl group because the methylene protons resonance was very broad in salt solu- tions and could not be read accurately. 8A Figure 12.--Chemical shift of DMSO protons XE: DMSO/NaAl(But)u with the baseline sub- tracted. . CPS 85 2- .Gln " CD C) ’\.r\ ,\ ,\ :QV‘KJ - “A’\IF I k’\’UTt i, I I I l I I I I I I I I I I I I II I 2 3 4 5 6 7 8 9 IO II I2 I3 I4>I5 I6 I7 I6 'DMSO R TIO _ A Nc", AL (BUII4 Figure 12 86 Figure l3.--A. Chemical shift of the N-methyl protons of 1M2PY vs. 1M2PY/NaA1(But)u mole ratio. NaA1(But)u concentration = 0.38 M, B. Blank. 87 GD ma themes no. ¢0. mm. mm. km. 00. mm. 0.... K. N5. 360 88 Figure 1A shows the behavior of the dioxane protons as a function of lM2PY/NaA1(But)u mole ratio. With the salt present, there is only a slight shift in the dioxane protons as shown in line A, Figure 1A. In order to determine what role the anion might have in producing a shift in the N-methyl protons reson- ance of 1M2PY, a large amount of Br- ion was added in the form of tetrahexylammonium bromide and the shift of the N-methyl resonance was plotted as a function of 1M2PY/Br” mole ratio. The results are shown in Figure 15. No bromide anion effect was observed. Admittedly Br' and A1(But)u- anions may produce different effects but Br- would be expected to be solvated much more than A1(But)u_ and thereby should produce a greater influence. The effect of anion cannot be completely ruled out, how- ever. Figure 15 illustrates that at low ratios of lM2PY/NaA1(But)u, the terminal methyl protons on the tetrabutylaluminate anion undergo a shift. The shift of the anion protons is very similar to that reported by Day. Figure 16 shows the nmr spectra of the terminal methyl protons and the two methylene group protons not bonded to the aluminum atom in the tetrabutylaluminate anion in solutions with DMSO/NaAl(But)u ratios of 0.25:1 and 18:1. Although there are some differences between the two spectra, there certainly are no extensive changes. The fact that there are no major changes in 89 Figure 1A.-—A. Chemical shift of dioxane protons gs. lM2PY/NaA1(But)u. NaA1(But)u concentra- tion = 0.36 M, B. Blank. 90 :Apsquaszstmsa m m a :H ohswflm _ _ 1, ,mam mam .mHN 8&0 91 Figure 15.--A. Chemical shift of N-methyl protons of 1M2PY vs, lM2PY/(Hex)uNBr mole ratio. Salt concentration = 0.38 M, B. Chemical shift of terminal methyl pro- tons of NaA1(But)u vs, lM2PY/NaAl(But)u mole ratio. NH Ha oapsm ofloz m m ma mhdwfim 92 j/ /I/ C) C) O C) C) 860 IJ/I mma 0 fl) :2 93 Figure 16.--NMR spectra of terminal methyl protons and methylene protons not bonded to A1 of tetrabutylaluminate anion in: A. DMSO/NaAl(But)u of 0.25/1 B. DMSO/NaA1(But)u of 18/1 Salt concentration = 0.36 M, 9A Terminal methyl protons Methylene protons not bonded to Al Figure 16 95 anion spectra could indicate that the anion is not highly solvated. This evidence supports the idea of minimum anion solvation. The proton resonance of the methylene protons bonded to the aluminum atom do exhibit some line broadening in going from a DMSO/salt ratio of 0.25:1 to 18:1 but the line is broad to begin with and is, there- fore, difficult to study. I Far Infrared Study of NaA1(But)u in Various Solvents Solutions of NaAl(But)u in DMSO and 1M2PY exhibit i cation solvent bands at 195 cm“1 and 206 cm”1 respectively. These energies correspond, within experimental error, to those found with more familiar salts such as sodium perchlorate, sodium iodide and sodium tetraphenylborate. The far infrared Na-DMSO cation~solvent band changes in appearance as the concentration of DMSO is decreased. Figure 17 shows the Spectra of NaA1(But)u in dioxane with ratios of DMSO/salt of 18:1 and 5.5:l. The spectra of NaA1(But)u in pure dioxane and pure heptane are also shown. The absence of strong absorption in the dioxane solutions in the 200 cm’1 region indicates a lack of strong solvent-ion interaction. The appearance of the Na-DMSO band in a mixed solvent system of DMSO and dioxane shows that sodium ion is undoubtedly solvated by the DMSO in the presence of dioxane. The Na-DMSO band occurs at about the same frequency in the presence of 96 Figure l7.—-Far infrared spectra: A. DMSO/NaA1(But)u ratio of 5.5:1 B. DMSO/NaAl(But)u ratio of 18 C. NaA1(But)u in dioxane D. NaAl(But)u in n-heptane Salt concentrations = 0.36 M, .. Rifle-«I. .. .! 97 150 200 250 O. O oocmpHOmp¢ Cm Figure 17 98 dioxane as the sodium bands in pure DMSO and with more normal salts such as the perchlorate. Even at the relatively low DMSO/NaA1(But)u ratio of 5.5:1, the Na-DMSO band is still evident. In pure dioxane and at very low DMSO/NaAl(But)u ratios the sodium solvent band in DMSO gives way to a much broader absorption. The broad absorption is also observed with heptane as sol— vent. The nature of the species which gives rise to the broad absorption has not been determined as yet. The absorption data in Figure 17 should not be taken as quantitative. A summary of the far infrared absorption studies in various pure and mixed solvents are tabulated in Table 12. Table l2.—-Infrared spectra of NaAl(But)u in {arious solvents from 320 cm-1 to 170 cm- . Q'm‘igfl NaAl(But)u _ Na-Solventl Solvent Concn. M v, max. (cm‘ ) DMSO 0.5 195 DMSO 18 0.5 197 NaAl(But)u = 1 QMéQ 212 0,5 200 NaAl(But)u 1 lM2PY 0.5 206 2PY 0.3 207 50:50 0.25 195 DMSO-Dioxane 50350 0.25 206 1M2PY-Dioxane Dioxane 0’8 broad absorption and 0.36 Heptane 0.A broad absorption V. MISCELLANEOUS OBSERVATIONS 1-Methy1—2—Pyrrolidone Complex with TiFu A solution of titanium tetraflouride and 1M2PY were mixed together in a dry box. A greyish white (probably impure) product precipitated. The infrared spectrum of the solid was obtained. The carbonyl band 1 frequency was lowered by 80 cm' , an extremely large shift. In the complex spectra, far infrared bands occurred at 625, 602, 553, A81, 35A and 285 cm_l. In pure solvent the bands were exhibited at 65A, 615, 557, A80, and 308 cm'l. It is probable that a complex was formed between 1M2PY and TiFu. Alkali Salt Solvation Studies in Other Solvents Tri-n-butylphosphate, the n-butyl ester of phos- phoric acid, and propylene carbonate are two solvents which have been studied in addition to the 2-pyrrolidones. The two compounds have structures listed below. Some of their physical constants are also listed. 99 100 H H2 0 HC-—-C-—-C III 3 I I (osut)3 0\ /0 C II 0 Tri-n-butylphosphate Propylene carbonate u 3.07 (76) u A.98 (77) e = 8.05 (79) E = 69 (78) M.W. 266.3 M.W. 102.09 The trialkylphosphates show numerous examples of interesting complex formation. Ferraro and co-workers (79) obtained far infrared spectra of complexes with the formula M(NO3)3:TBP where M is a rare earth metal. Iodine forms a complex with tributylphosphate which re- sults in the visible iodine band being shifted from 525 mu to A50 mu in n-heptane solution (80). Conven— tional thermodynamic constants were evaluated by study— ing the equilibrium as a function of temperature. One of the most common uses of tributylphosphate is as an extractant. For example, lithium halides can be extracted from aqueous solution by tributylphosphate. In distribution studies Morris and Short found that lithium maintained a hydration number of A21 in the organic phase (81). The phosphoryl oxygen atom has been shown to be the sight of interaction in complexes such as FeCl3zOP(OEt)3 (82). lOl Propylene carbonate has a narrow but useful spectral window in the far infrared region at a frequency that is appropriate for the study of lithium solvent inter- actions. The high dipole moment and high dielectric constant make it suitable for the investigation of ion-dipole interactions. The alkali ion vibrations which have been observed in these solvents are listed in Table 13. These two solvents are not as useful as the pyrrolidones because they are not sufficiently transparent and solubilities of alkali salts are low. One of the most interesting observations with tri- n—butylphosphate is that the alkali halides decompose the solvent. A white crystalline solid was filtered from the reaction mixture of Lil and TBP. The supernatant liquid gave a negative test for iodide when tested with silver nitrate which indicates that one of the reaction products is butyliodide. The ir spectrum of the solid matched the literature spectrum for LiPOu(But)2 (83). The carbon—hydrogen—phosphorus analysis is given below and agrees well with that of LiPOu(But)2. Carbon Hydrogen Phosphorus Observed % AA.O6 8.18 1A.09 Calculated AA.A5 8.39 1A.33 Qualitatively, the rate of decomposition depends on the halide anion with I>Br>Cl. 102 TABLE l3.——Absorption spectra of alkali metal and ammonium ions in prOpylene carbonate and tri-n-butylphosphate. Vmax (cm-l) Salt Propylene carbonate Tri-n—butylphosphate LiClOu 386 i 6 cm-1 38A i 6 cm-1 Lil 388 " LiBr 380 " NaBF,4 --— " 208 " NaClOua 203, 167 " 208 " NaSCN 216, 167 " 219 " NHuClOu 203, 167 " 218, (233)b " NHMSCN 203, 167 " 218, (233)b " aOverlapping bands of approximately equal intensity appear in the sodium and ammonium solution Spectra in prOpylene carbonate. bA shoulder is present at 233 cm- spectra. 1 in these Table 1A lists several other solvents which were briefly examined in the 660 to 330 cm_1 region. All of these solvents suffer from either poor transparency or low salt solubility. Low Temperature Spectra of Acetone Solutions of Sodium Iodide and Ammonium Thiocyanate The spectra of acetone solutions of 0.1 M sodium iodide and 0.3 M ammonium thiocyanate were examined at 103 —70°C and compared to the spectra at 30°C. No signifi— cant differences were observed in intensity or absorbance over the one hundred degree range. The problem is com- plicated by the fact that the cation solution spectra display several overlapping bands so that small changes could easily go unnoticed. TABLE 1A.—-Solvents which were eliminated for far infrared studies. Formamide y-thiobutryolactone Dimethylacetamide l-methy1-2—pyridone N—methylacetamide 1—ethyl-2-pyrrolidone Hexamethylphosphoramide l—buty1-2~pyrrolidone l,l,3,3,-tetramethy1urea Isopropanol Benzonitrile Methanol Thiophosphoryl chloride Tetramethylenesulfoxide VI. SUGGESTIONS FOR FURTHER STUDY 1. Solutions of divalent and trivalent cations could be investigated for the possible existence of ion vibrations. The higher charges would tend to produce a stronger ion—solvent bond but the large masses would favor a low frequency ion—solvent vibration. These ions are more amenable to coordination number determinations by nmr techniques. It is very possible that Raman spectroscopy would be useful to characterize the cation- solvent symmetry if an actual complex existed in solu- tion. 2. Another test of the alkali ion vibration for anion independence would be to test the molar absorp— tivities of as many salts of a given cation as possible. If the vibrations are truly independent of the anion, then all the molar absorptivities of the different salts should be identical. If this experiment is performed the salts must all be anhydrous and the experiment should be performed as described in the Beer's law experiment with LiClOu in the Results and Discussion section. If a suitable solvent is found, integrated intensities should be measured. 10A 105 3. Conductance data, including concentrations useable in the far infrared work would provide additional information about ionic association. Perhaps the con- centration of ion pairs at high (1 M) salt concentrations could be calculated and related to the infrared bands. A. Some low temperature far infrared work has already been described. Low temperature work is severely limited by solvent freezing points. Many of the solvents studied have boiling points above 200° at one atmosphere. High temperature studies appear to be more feasible than low temperature studies. If a dependence of the far infrared ion vibrations on temperature could be found, information on the energy of the ion-solvent interaction could be obtained. 5. Water can be studied in the far infrared (8A). Low energy (~A00 cm-l) divalent and trivalent cation- water Raman bands have been reported (85). It is likely that some of the Raman bands are also infrared active. It would be interesting to study these bands with water as solvent. BIBLIOGRAPHY 10. ll. 12. 13. 1A. 15. BIBLIOGRAPHY W. Maxey and A. I. Popov, J. Amer. Chem. Soc., 89. 2230 (1967). W. Maxey and A. I. Popov, Ibid., 28, AA70 (1968). W. Maxey and A. I. Popov, Ibid., 91, 20 (1969). . F. Edgell, A. T. Watts, J. L. Lyford and W. M. Risen, J. Amer. Chem. Soc., 88, 1815 (1966). . F. Edgell, Abstracts, 153rd National Meeting of the American Chemical Society, Miami Beach, Fla., April 9-1A, 1967, No. R—1A9. . F. Edgell, J. Lyford and J. Fisher, Abstracts, 155th National Meeting of the American Chemical Society, San Francisco, Calif., March 31— April 5, 1968, No. S-136. . J. French and J. L. Wood, J. Chem. Phys., 82, 2358 (1968). . D. Dyke, P. G. Sears and A. I. POpov, J. Chem. Phys., 7;, A1A0 (1967). Olavi I. Virtanen, Suomen Kemistilehti, B, 32, 257 (1966). . E. Zaugg, J. Amer. Chem. Soc., 89, 2903 (1960. C. 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D. Smithson and A. L. Rozek, "Infrared Spectra and Characteristic Frequencies 700- 300 cm- ," Interscience Publishers, New York, N.Y., 1968. E. D. Palik, J. Opt. Soc. Amer., 88, 1329 (1960). J. Stewart in "Interpretative Spectroscopy," S. K. Freeman, ed., Reinhold Pub. Corp., New York, N.Y., 1965, pp. 131—169. J. W. Brasch, Y. Mikawa and R. J. Jakobsen, App. Spect. Rev., 3, 187 (1968). J. R. Ferraro, Anal. Chem., A, 2AA (1968). E. V. Loewenstein, Appl. th., 3, 8A5 (1966). C. C. Helms, H. W. Jones, A. J. Russo and E. H. Siegler, Jr., 8pectrochim. Acta, 33, 819 (1963). N. T. McDevitt and A. D. Davidson, J. Opt. Soc. Amer., 88, 1695 (1965). E. Schaschel and M. C. Day, J. Amer. Chem. Soc., 38, 503 (1968). T. P. Whaley in "Inorganic Synthesis," T. Moeller, ed., McGraw—Hill Book Co., Inc. New York, N.Y., vol. v, 1957, p. 6. D. A. Skoog, D. M. West, "Fundamentals of Analytical Chemistry," Holt, Rinehart and Winston Co., New York, N.Y., 1956, Chapter 12. K. G. Stone, "Determination of Organic Compounds," McGraw—Hill Book Company Inc., New York, N.Y., 1956, p. 6. 110 51. J. E. Katon, W. R. Feairheller, Jr., and J. V. Pustinger, Jr., Anal. Chem., 38, 2126 (196A). 52. K. Nakamoto, "Infrared Spectra of Inorganic and Coordination Compounds," John Wiley and Sons, Inc., New York, N.Y., 1963, p. 107. 53. S. S. Mitra and P. J. Gielisse, "Infrared Spectra of Crystals," AFCRL-69-393, June (1965). 5A. C. M. Randall, R. M. Fuller and D. J. Montgomery, Solid State Comm., 3, 273 (196A). 55. A. J. Parker, Quart. Rev., 88, 163 (1962). 56. F. J. Rossotti and H. Rossotti, "The Determination of Stability Constants in Solution," McGraw— Hill Book Company, Inc., New York, N.Y., 1961. 57. G. Donnay and J. W. Gryder, J. Chem. Ed., fig, 223 (1965). 58. H.G.F., Winkler, Acta Cryst., 6, 99 (1953). 59. A. A. Maryott and F. A. Smith, "Table of Dielectric Constants of Pure Liquids," NBS Circular 51A, Aug. 10, 1951. 60. R. Sangwald and A. Weissberger, Physik. Z., 38, 268 (1929). 61. R. M. Bedger, J. Chem. Phys., 8, 128 (193A). 62. R. M. Badger, Ibid., 3, 710 (1935). 63. C.H.D. Clark, Phil. Mag., 18, A59 (193A). 6A. H. S. Allen and A. K. Longair, Nature, 135, 76A (1935). 65. F. A. Cotton and G. Wilkinson, "Advnaced Inorganic Chemistry," Interscience Publishers, 1962, p. 2A9. 66. H. H. Willard, L. L. Merritt, and J. A. Dean, "Instrumental Methods of Analysis," Ath ed., D. Van Nostrand Co., Inc., Princeton, New Jersey, 1965, p. 12A. 67. R. E. Hester and R. A. Plane, J. Chem. Phys., 88, A588 (1966). 68. 69. 70. 71. 72. 73. 7A. 75. 76. 77. 78. 79. 80. 81. 82. 83. 8A. 85. . E. Hester and K. Krishnan, Ibid. . Davidson, Trans. N.Y. Acad. Sci 111 9 £9, 3A05 (1967). E. Hester and K. Krishnan, Ibid., A7 (17A7 (1967). .. 29. 316 (1958). P. Smyth, W. S. Walls, J. Chem. Phys., 3, 557 (1935). . A. Cotton and G. Wilkinson, "Advanced Inorganic Cheméstry," Interscience Publishers, 1962, p. l 0. . J. Orville-Thomas, Chem. Revs., 57, 1179 (1957). . Fratiello and D. P. Miller, Mol. Phys., 8;, 37 (1966). S. Frankel, T. R. Stengle and C. H. Langford, Can. J. Chem., 88, 3183 (1968). K. Estok and W. W. Wendlandt. J. Amer. Chem. Soc., 11. A767 (1955). L. McClellan, "Tables of Experimental Dipole Moments," Freeman, San Francisco, 1963. Gutmann, Private communication. R. Ferraro, C. Cristillahi and 1. Fox, J. Inorg. Nucl. Chem., submitted for publication. . Tsubomura and J. M. Kliegman, J. Amer. Chem. Soc., 82. 131A (1960). C. Morriss and E. L. Short, J. Inorg, Nucl. Chem., 32: 291 (1963). . R. Bronstein and M. A. Bredig, J. Amer. Chem. 888,, 88, 2077 (1958). D. Smith, J. Inorg. Nucl. Chem., 9, 150 (1959). . A. Draegert and D. Williams, J. Chem. Phys., A8, A01 (1968). . A. Plane, Cornell University, private communica- tion, 1969. 86. 87. 88. 89. 112 . T. McDevitt, A. L. Rozk and A. D. Davidson, Technical Documentary Report ML TDR 6A-l92, Air Force Materials Laboratory, Wright- Patterson Air Force Base, Ohio, 196A. Yoshinaga, S. Fujita, S. Minami, A. Mitsuishi, R. A. Oetjen and Y. Yamada, J. Op. Soc. Amer., A8, 315 (1958). Herzberg, "Infrared and Raman Spectra," D. Van Nostrand Co., Inc., Princeton, New Jersey, 196A, p. 182. Herzberg, Ibid., p. 172. APPENDICES APPENDIX I COMPUTER PROGRAM FOR SPECTROPHOTOMETER CALIBRATION 115 The calibration data accompanying the 301 spectro- photometer should be valid indefinitely unless the grating adjustments are tampered with or accidentally altered. At one time, all the gratings had to be recalibrated because of a loose grating mount. In calibrating the instrument, it was found that the most useful and detailed source of calibration infor- mation for the 301 spectrophotometer has been compiled by McDevitt (86). Exact operating conditions have been listed by him and will not be discussed here. Useful calibration data is also contained in the 301 manual. Yoshinaga (87) published information for the region below 50 cm-l. A plot of wavenumber vs. drum number is linear for the 301. Instead of making the plot by hand, the data can be processed with the aid of the program shown in Table 15. Sample cards are shown in Figure 18. C) 116 TABLE l5.—-Program Cal. PROGRAW CAL DIMENSION 1(100). Y(100) 1 FORMAT (12V. I2. 6!. 11. 5X. 11. 5X. Ii) 2 FORMAT (I3. 2X. 514.8, 5X, E14.8) 3 FORMAT (1X, 12. 6X, F8.2. 2X. F9.4. 1X. F10.5. 10X. F10.6) 10 FORMAT (62H I X YIOBS) Y(CALC) YIOBSI- 5YICALC)) 16 FORMAT (1H1 35x. 12. 20H LINE PER MM GRATING. 3x.12./7/ 9x." 13380809 N0. FREQ CM-i DIFF.. /I * 18 FOQMAT (1X) - 19 FOAMAT I9x. r7.1. 4x2 F10.5, 4x. r9.5. so FOQMAT (1H1 35x. 12. 208 LINE PER MM GRATIMG. 3x.12./// 9x. 129AFnEo CM-i DRUM yo. DIFF.) /) 55 FORMAT (9X, F7.2. 3x. F8.2. 4x. F6.2) 39 FUQMAT(1H1. 2x. 12. 90H LIVE pea MM GRATING) 45 FOQMAT (1H8, 13. 2x, E14.8. 5x. E14.8) 42 RtAD 1. J. L1. L2. L3 IF (J) 40. 40. 41 41 PRINT :9. J NAME = J IF (Li-2) 4. 4. 12 CALCULATE LEAST SQUARES COASTANTS 4 SO : 0,0 5X = 0.0 SY = 0.0 SXY : 0,0 I31 5 READ 2. J. X(I). YIII IF (J'I) 73616 7 SU=SO+X(I)¢*2 SX=SX+Y(I) SY=SY+Y(I) SXV=SXV+X(I)*Y(I) I=I+1 GO TO 5 5 131.1 A ' (SX*SXY-SYtSO)/(SX**2 ' 50*FLOATF (II) B = (Sv-sx -SXY*FL0ATF(I))/(SX**2-SO*FLOATF(I)I PRINT 45. I. A, R It (Li-2) 11.8.11 8 PRINT 10 DUO K:1. I YCALC:A+B¢Y(K) DIFF=Y(K)-YCALC 9 PRINT 3. K. X(K). Y(K). YCALC. DIFF 11 PRINT 18 PRINT 18 pRINT 45- I: A: a 12 II (L2-2) 25. 25. 13 CREATE Y(X) TABLE 13 IF (Ll-2) 15. 15. 14 14 RtAnz. M. A, B 15 READ 2. M, x0, 0* Y0 = 0.0 NPzM/5n DO 20 IP:1,NP 117 TABLE l5.--Co§tinued. 17 20 54 £4 25 U 26 27 28 29 61 62 56 67 46 40 RU~110300 IF (”P-1°) s4. pRINI 16. NAME. 00 20 LN=1.50 VCALC : t+flt¥0 DIFF=Y7A;COYO PRINT 19. XI). YUsYCAMC XU=XC+PX Ip=Ip+I NLQ: M—5CoNP PHINI 16. NAME. DU 24 qV:1, MLR YCALC=A+3ch DIFFzYCALC-YO pRINT 19. X0. YU=YtAgc XU: X060! PRINT 18 PRINT 18 PRINT 45. M, A, R II (L3-2) 43. 43: CK5ATE X(Y) TABLE 1+ ((1.2) 79, 99. IF (L2-2) 78. 28. k'IIAI.) 2. v. A. q Akin}, 4. v0. 07 AP:.A/2 BP:1,U/9 NP:M/50 X0 :0,” 00 32 79:1. IF (AP-I3) PRINT 50. NAMF. DU 32 1V=10 SO XCALC : AP+BPoYJ DIFF=XTALC-XC PRINT ‘8. V0. x0=XCAfc YO :VO+DV IP:IP+1 MLQ : v-30omp PRINT 10. NAME. 00 37 gNzl. NLQ XCALC : AP+BPtYO DIFF:X7ALCeX0 PRINT ‘8. Y0. XU:XCALC YU:YF+DY PRINT 18 PRINT 18 pRINT 45. M! GO TO 42 STWP EN” 17.17 IP YCALC. DIFF ID YCALC: DIFF 96 27 99 we 36, 31, 31 IP XCALCI DIFF Ip XCALC. DIFF A9. 88 20007 118 Figure 18.—- Sample data cards for program cal. 1. Identification card; Run number 20. 2. Data Card; The drum number (xobs = 209M.O) with corresponding wave number (Yobs = 335.3“) are entered. 3. This card is the same for all sets of data. M. The number of lines of calculated wavenumber (230), a starting drum number (0) and the drum number incre- ment (10) are entered. 5. The number of lines of calculated drum numbers (230), the starting drum number (115), and the wave number increment (1) are entered. 6. Blank. 119 6 999999999999999900099330999990099090 830 [315367 9 9'0Hunl‘lsnh'n"n.122;9:.6””2929””nnu153 5 9:39:92999.999900900900990999,09999009900 9 IOll"IIIJNISIIIHIIONIIZID I 3‘ ) QC 230 00050000 TEE”, m” M m 4 99-999L ...... .,00009909999990,0,99999999090000 l1 3 C 9 9 7 I 910Hl2UM”IIIIIIIINflRDNaflZTRRDJI12133433113190‘120)“t 90.09.00.00000000000000000009000000000000000000OI I 2 J C 9 9 7 I ll." ummsnnunannnuannaann12131433323304 czuuauuaa 3094.0 335.34 0000000.000,000000000000000000000000000000000000000000 I 3 J 4 9 I T 9 I IonIzumsunnumvunaaannann123334«3173304 ‘ZCJMOS$¢IC0$SISIJJ¥ RUN NUMBER 20 2 3 000.0000000.00000000990000000990000000000000000000000000T 345:7.omununnununanannaannnnnnnuwxnanouuuuuwnanmmnuuusnau lllllllllllllllllllllllllllllllllllllllllllllIlllllllllll 2222222.2222.2222222.22222222222222222222222222222222222222 33333333333333333333333333.33333.33333333333333333333333333 4 444:.444444444444444444444444444444444444444444444444444! 55-5 555.55555555555555555955555555555555555555555555555555 66666566666666566668666665666556865566655666666666666666566 1777777777777777777777777777777777777777777777777777777777? 8.888888888888888888888888888888888888888888888888888888886 9999999 99999999999 99999999999999999999999999999 I . :0 ;' -: ,. a :5 ma ,7 's 13 .J r: 2: :; 2; r ;a " PF :9 ;c s .; A- .: . ' ,; in .. ar ;‘ 4‘ 4' a: s. It a: :4 ‘J .3 5: a? .3 5: :3 3c :‘ 25' 99999999 2305579° Cg "‘ Figure 18 APPENDIX II SAMPLE FREQUENCY CALCULATIONS APPENDIX II For a tetrahedral XYu model based on valence forces the v3 and vu triply degenerate vibrations are related by the following formula (88): Am A A = §_E_EL (1 +.__X ) 3 u 2 2 m 1 mx y where v . l 3 = 13' /2nc in cm- , ‘Au' /2flc in cm-l, vu = k = force constant in dynes/cm of the XY bond, k' = force constant in dynes/cm of the YXY angle deformation, 1 = XY distance in cm, m = mass of atom Y m = mass of atom X c = speed of light in cm/sec For our purposes Y is considered to be a solvent molecule and X is considered to be a cation. If the low energy angle deformation VA and the Quantity EEK, are assumed constant in going from 12 121 122 + Li7(1M2PY)u to Li6(lM2PY)u+, the following equation gives an estimate of the expected frequency of the Li6(lM2PY)u+ v3 vibrations. Am A (l + 1M2PY )l/d v 6 le V 7 3L1 Mm . (1 + __E9321 )1/2 mLi7 The L17—1M2PY infrared band occurs at 398 cm'l. The 6 calculated value for the Li —1M2PY band is 427 cm“ 1 Similar calculations can be performed using a diatomic (62) or triatomic model (89). Li7-1M2PY to L16 In going from -1M2PY the calculated values are H26 cm- and 427 cm”1 for the linear diatomic and linear tri- atomic models, respectively. 1 w, \llllWlHlllllUlH l 5 4| 6 5 8 7 1 3 0 3 9 2 4| “I'll "I “ E "I u 9 u 9 “I