SPECTROSCOPIC STUDIES OF COMPLEX COMPOUNDS IN NONAQUEOUS SOLVENTS Thesis for the Degree of Ph. D. MICHIGAN STATE UNIVERSITY WILLIAM JAN MCKINNEY 1 96 9 " LIE‘RA' Y Michigan .; are University THES'S This is to certify that the thesis entitled SPECTROSCOPIC STUDIES OF COMPLEX COMPOUNDS IN NONAQUEOUS SOLVENTS presented by William Jan McKinney has been accepted towards fulfillment of the requirements for Ph .0. degree in ChemlStry V 6/50 flMéJESD/gm/ Major professé/ Date July 29, 1969 0-169 BIND!“ IY BUAB a SUNS' I. ,3; ; er VA’Y INC. I’fi'fr‘s ‘ SPECTROSCOPIC STUDIES OF COMPLEX COMPOUNDS IN NONAQUEOUS SOLVENTS BY William Jan McKinney AN ABSTRACT OF A THESIS Submitted to Michigan State University in partial fulfillment of the requirements for the degree of DOCTOR OF PHILOSOPHY Department of Chemistry 1969 ABSTRACT SPECTROSCOPIC STUDIES OF COMPLEX COMPOUNDS IN NONAQUEOUS SOLVENTS PART I: PYRIDINE-IODINE COMPLEXES PART II: ALKALI METAL IONS IN PYRIDINE AND ACETONE SOLUTIONS BY William Jan McKinney Spectroscopic techniques have been used in this investigation to study the pyridine-iodine complex as well as the solvation of alkali metal salts in pyridine and acetone. Pyridine and its homologues are known to form charge-transfer complexes with iodine, the complex forma— tion constant being dependent on several physical and chemical parameters of the system. The effect of sub- stitution on the pyridine ring on the formation constants of the complexes has been determined spectrOphotometrically in carbon tetrachloride solutions at 25°. A plot of log K (t mole-1) gg, the Hammett o constants of the pyridines M yields a straight line given by the equation tog KM = -2.25 o + 2.11, indicating that the addition of nucleOphilic William Jan McKinney substituents to the pyridine ring increases the strength of the complex. Likewise a plot of log K X§° pKa of the M respective pyridines follows a linear relationship with the exception of cases in which steric hindrance becomes an important factor. If the substituent groups are elec- tron donors themselves (phenyl or nitrile), a simple spectrophotometric technique fails to give the formation constant of the complex, presumably owing to the stepwise formation of two complexes. Formation constants in mole fraction units of the pyridine-iodine charge-transfer complex have been deter- mined at 25° in twelve solvents with dielectric constants varying from 1.92 to 10.36. The Kx values range from 612 in nfhexadecane to 3248 in g—dichlorobenzene. The enthalpies and entropies of the complex formation have been determined in five solvents. Comparison of the data with the solvent properties indicates that an increase in the dielectric constant of the reaction medium leads to the stabilization of the pyridine-iodine complex. The phenomenon is complicated, in some cases, by specific solute-solvent interactions, such as solvation of pyridine by chloroform or by the formation of the triiodide ion in polar solvents. The solvation of lithium, sodium, and ammonium salts by pyridine was studied by infrared techniques. A new band was observed in these solutions which was William Jan McKinney attributed to a cation-solvent vibration. Studies of the lithium- and ammonium-solvent bands by isotopic substi- tution techniques tend to confirm the assignment of this band to a cation-solvent vibration. The assignment is also supported by the splittings of several pyridine skeletal vibrations in alkali metal salt solutions. The magnitude of these splittings shows that the strength of the solva- tion of the above cations by pyridine increases in the order Na+ < NH: < Li+. The observed cation—solvent bands show little anion effect, except for the sodium iodide salt. However, a study of the spectra of the perchlorate ion shows that the symmetry of the ion is lowered in pyridine solutions of the lithium, sodium, and ammonium salts. Some preliminary work in the far-infrared region of the spectrum has been done in acetone solutions. Solvent-cation bands for lithium and sodium salts have been observed in this region. The bands for the lithium salts appear to be anion dependent, especially the lithium chloride and lithium bromide salts. Isotopic studies with the lithium salts and d6—acetone tend to support the assign- ment of the bands in the far-infrared to cation—solvent vibrations, although lithium salts in d6-acetone give ambiguous results. SPECTROSCOPIC STUDIES OF COMPLEX COMPOUNDS IN NONAQUEOUS SOLVENTS BY William Jan McKinney A THESIS Submitted to Michigan State University in partial fulfillment of the requirements for the degree of DOCTOR OF PHILOSOPHY Department of Chemistry 1969 ACKNOWLE DGMENTS Without the patient guidance and encouragement of the major professor for this paper, Dr. Alexander I. Popov, the work could not have been completed. The author also gratefully acknowledges the support of this work by grants from the National Science Foundation and the 0.8. Army Research Office (Durham). Special appreciation goes to my colleague, Ronald H. Erlich, who spent many hours of his time proofreading the manuscript. He and others in this laboratory--among them Dr. Frank M. D'Itri, Ming Keong Wong, Delores Bowers, and Dr. Richard Nicholson--have contributed to the work through their support, encouragement, suggestions, and diversionary activities. Finally, the author gratefully acknowledges his wife, Kathy, for typing the manuscript as well as her emotional support. ii TABLE OF CONTENTS ACKNOWLEDGMENTS O O O O O O O 0 C C C C O C O 0 LIST OF TABLES O O O O O O O O O O O O O C O O 0 LIST OF FIGURES . C O O O O I I O O O I O O O 0 PART I. PYRIDINE-IODINE COMPLEXES . . . . . . . PART INTRODUCTION . . . . . . . . . . . . . . . HISTORICAL PART . . . . . . . . . . . . . . Acceptor Strength . . . . . . . . . . . . Steric Considerations . . . . . . . . . . Donor Strength . . . . . . . . . . . . . Solvent Effects . . . . . . . . . . . . . EXPERIMENTAL . . . . . . . . . . . . . . . Chemicals . . . . . . . . . . . . . . . . Preparation and Stability of Solu ions . Spectral Measurements . . . . . . . . . . Calculations . . . . . . . . . . . . . . RESULTS AND DISCUSSION . . . . . . . . . . Donor Strength and Steric Considerations Solvent Effects . . . . . . . . . . . . . II. ALKALI METAL IONS IN PYRIDINE AND ACETONE SOLUTIONS . . . . . . . . HISTORICAL INTRODUCTION . . . . . . . . . . Solvation . . . . . . . . . . . . . . . . Pyridine O O O O O O O O O O O O O O O 0 Acetone . . . . . . . . . . . . . . . . . EXPERIMENTAL . . . . . . . . . . . . . . . Chemicals . . . . . . . . . . . . . . . . Preparation of Solution . . . . . . . . Instrumentation . . . . . . . . . . . . . Experimental Techniques . . . . . . . . . iii Page ii viii 80 81 81 82 85 88 88 9O 9O 91 Page RESULTS AND DISCUSSION . . . . . . . . . . . 92 Infrared Spectra of Alkali Metal Salt Solutions in Pyridine . . . . . . 92 Infrared Spectra of Alkali Metal Salt in Acetone and d6-Acetone . . . . . . . . 100 MDENDUM I O O O C O O O O O O O O O O O O O O O O 104 LITERATURE CITED . . . . . . . . . . . . . . . . . 107 APPENDIX 1. COMPUTER PROGRAM OF THE KETELAAR EQUATION FOR THE CALCULATION OF FORMATION CONSTANTS . . . . . . 115 APPENDIX 2. CALCULATION OF THE EFFECT OF A COMPETING EQUILIBRIUM WITH THE SOLVENT ON THE FORMATION CONSTANT OF A PYRIDINE-IODINE COMPLEX . . . 119 APPENDIX 3. SUGGESTIONS FOR FUTURE WORK . . . . . 123 iv Table 1. ' 10. ll. 12. 13. 14. LIST OF TABLES Literature Values for the Formation Constant of the Pyridine-Iodine Complex in Several Solvents . . . - . . . . . . . . Physical Constants of Several Substituted Pyridines . . . . . . . . . . . Experimental Data on the 2—Fluoropyridine- Iodine Complex in Carbon Tetrachloride . . . Experimental Data on the 2,5-Lutidine- Iodine Complex in Carbon Tetrachloride . . . Experimental Data on the 3-Chlor0pyridine- Iodine Complex in Carbon Tetrachloride . . . Experimental Data on the 2—Chloropyridine- Iodine Complex in Carbon Tetrachloride . . . Experimental Data on the 2-Bromopyridine- Iodine Complex in Carbon Tetrachloride . . . Experimental Data on the 4-Ethylpyridine- Iodine Complex in Carbon Tetrachloride . . . Experimental Data on the 3-Bromopyridine— Iodine Complex in Carbon Tetrachloride . . . Experimental Data on the Pyridine-Iodine Complex in Carbon Tetrachloride . . . . . . Experimental Data on the 4-25Butylpyridine- Iodine Complex in Carbon Tetrachloride . . . Experimental Data on the 4-Picoline-Iodine Complex in Carbon Tetrachloride . . . . . . Experimental Data on the 3-Picoline-Iodine Complex in Carbon Tetrachloride . . . . . . Experimental Data on the 3,5-Lutidine- Iodine Complex in Carbon Tetrachloride . . . Page 11 16 29 30 31 32 33 34 35 36 37 38 39 4O Table Page 15. Experimental Data on the 3,4-Lutidine- Iodine Complex in Carbon Tetrachloride . . . 41 16. Experimental Data on the 3-Ethylpyridine- Iodine Complex in Carbon Tetrachloride . . . 42 17. Experimental Data on the Pyridine-Iodine Complex in Carbon Tetrachloride . . . . . . 43 18. Experimental Data on the Pyridine-Iodine Complex in Dichloromethane . . . . . . . . . 46 19. Experimental Data on the Pyridine-Iodine Complex in 1,2-Dichloroethane . . . . . . . 48 20. Experimental Data on the Pyridine-Iodine Complex in Chloroform . . . . . . . . . . . 50 21. Experimental Data on the Pyridine-Iodine Complex in B-fleptane . . . . . . . . . . . . 52 22. Experimental Data on the Pyridine-Iodine Complex in 1,1,2-Trifluorotrichloroethane . 54 23. Experimental Data on the Pyridine-Iodine Complex in meichlorobenzene . . . . . . . . 55 24. Experimental Data on the Pyridine-Iodine Complex in anexadecane . . . . . . . . . . 56 25. Experimental Data on the Pyridine-Iodine Complex in g—Dichlorobenzene . . . . . . . . 57 26. Experimental Data on the Pyridine-Iodine Complex in peXylene . . . . . . . . . . . . 58 27. Experimental Data on the Pyridine- Iodine Complex in Benzene . . . . . . . . . 59 28. Experimental Data on the Pyridine- Iodine Complex in Toluene . . . . . . . . . 60 29. Formation Constants (lmole-l) of Iodine Complexes with Pyridine and Substituted Pyridines in Carbon Tetrachloride Solutions at 2 25° . . . . . . . . . . . . . 62 30. Fonmation Constants of the Pyridine- Iodine Complex in Different Solvents . . . . 63 vi Table 31. 32. 33. 34. 35. Enthalpy and Entropy Values for the Pyridine-Iodine Complexation in Different Solvents . . . . Splitting of Three Pyridine Skeletal Vibrations by Li+, Na+, and NH+ Ions Absorption Bands of Alkali Metal Salts in Pyridine . Position of Several Perchlorate Bands in Pyridine Solutions of the Tetrabutyammonium, Lithium, Ammonium, and Sodium Salts Absorption Bands of Alkali Metal Salts in Acetone and d6-Acetone Vii 4 Page 79 93 95 99 101 LIST OF FIGURES Figure Page 1. Absorption spectra of the pyridine-iodine system in carbon tetrachloride solutions . . 22 2. Relationship between log KM and Hammett a function for the pyridine-iodine complexes in carbon tetrachloride solutions . . . . . 66 3. Relationship between the acidity constant of the pyridines in aqueous solutions and the formation constants of the pyridine— iodine complexes in carbon tetrachloride solutions . . . . . . . . . . . . . . . . . 69 4. Relationship between the formation constant expressed in mole fraction units of the pyridine-iodine complex and the dielectric constant of the solvent in which it was measured . . . . . . . . . . . . . . . . . . 72 5. Relationship between the wavelength of maximum absorption of the blue-shifted iodine band of the complex and the dielectric constant of the solvent in which it was measured . . . . . . . . . . . 76 viii I. PYRIDINE-IODINE COMPLEXES INTRODUCTION The beginning of interest in molecular complexes probably dates back to 1927 when Pfeiffer published a classic review, "Organische Molekiilverbindung."l Instances of additive combination between aromatic hydrocarbons and other organic compounds as well as certain inorganic com- pounds were known at that time but were in apparent viola- tion of the existing rules of chemical bonding. The sta- bility of these additive compounds varied a great deal. Some formed stable solid adducts exhibiting integral stoichiometries while the existence of others could only be inferred from changes in color or other physical prop- erties when the reactants were mixed together in solution. Pfeiffer explained the bonding between the components by postulating the existence of secondary valence forces within aromatic nuclei which were susceptible to satura- tion through interactions with quinones and various other molecules. After Pfeiffer's article a basis for understanding these compounds in terms of acid-base theory was given by Lewis.2 Interest in these compounds (molecular complexes) was stimulated by the discovery by Benesi and Hildebrand of a new absorption band in the ultra-violet region of the spectrum in solutions of benzene and iodine which was characteristic of a complex.3 This band provided a new means of studying the benzene-iodine complex and similar complexes, and its interpretation lead Mulliken to an extension of the Lewis acid-base theory in a quantum- theoretical form which provided the basis for the inter- pretation of a wide variety of phenomena associated with the molecular complexes.4 Since that time there has been a large increase in the literature available on the subject of molecular charge-transfer complexes.5 The interest in charge-‘ transfer complexes has been largely stimulated by the rich variety of possible complexes, their importance in biological systems, their importance as reaction inter- mediates and, to some extent, by the relative ease of experimental studies. As a result, there has been a large number of studies aimed at gaining an understanding of the factors affecting the stability of these species in solution. These factors can be thought of as belonging to one of four broad categories: 1. acceptor strength 2. donor strength 3. steric considerations 4 . solvent effects. Thus it was of interest to study a homologous series of complexes in order to determine the effect of each of the above factors on the stability of molecular complexes. The pyridine-iodine complex and its homologues were chosen for this purpose. HISTORICAL PART The existence of a solid 1:1 molecular addition compound of iodine with pyridine (Py) was first reported 6 Chatelet also reported the isolation by Chatelet in 1933. of two hydrated complexes which he identified as IZ-Py(H20)6 and IZ-Py4(H20)24. Most of the evidence for the formation of addition compounds of iodine with organic bases comes from spectral studies. Zingaro, e£_gl., have investigated the behavior of iodine in pyridine solutions spectrophotometrically and postulated the reaction to be: 4..._ + - 212 + Py ~—v IPy + I3 Infrared studies of iodine with pyridine and 2- picoline in carbon tetrachloride by Glusker, et_gl.8 showed a definite shift of the absorption band from those of the individual compounds upon the addition of iodine to the respective organic base. The shift of absorption bands was attributed to the formation of a halogen-amine com- plex. Mulliken and Reid,9 in their study of the pyridine- iodine system in heptane solutions, reported the existence of the 1:1 molecular complex. Upon the addition of pyridine, the 520 mu peak of iodine in heptane was shifted to 422 mu. This peak was attributed to the absorption by the complex, and the association constant for the complex was calculated to be 290 at 16.7°. Mullikeng'4b suggested the existence of two types of complex for the pyridine-iodine system: an "outer complex" (A) and an "inner complex" (B) with the S tructures ,1 I . + (UN: I (JN- < - , \I < I (A) (B) Kortfim and Wilski prOposed the same possibility on the basis of conductivity measurements of the pyridine-iodine system.10 11 isolated two different addi- Glusker and Miller tion compounds formed by 4-picoline and iodine: compound II has an ionic structure and compound I is a molecular addition analogous to the addition compound of pyridine and iodine prepared by many workers. Their studies have led to the postulation that compound I is an intermediate in the formation of compound II, and the mechanism of the reaction is as follows: CH CH H ._ 3 R. 3 it; , 3 6 11+ 1: :1 f) .—— + I I“. 1‘ 5‘9 '39 31:89 T— I 9 CH3 CH3 1 + Q (:4 “3C ON-I N69 N H 1‘ I C) C‘ 3 12 Crystallographic studies by Hassel, e£_21., on the 1:1 addition compound of 4-picoline and iodine showed that the I—I-N arrangement in compound I is linear with the I—I distance of 2.83 A0 while the I—N distance is 2.31 A0. Investigation by Hassel and Hope13 on the reac- tion product of pyridine and iodine prompted them to postulate that the cation of compound II has a linear structure: + ON - —- © H ETC 1 C I instead of addition next to the N atom on the Py ring as proposed by Glusker and Miller.11 Acceptor Strength A series of complexes between the Lewis bases; pyridine, 2-picoline, and 2,6-lutidine; and the Lewis acids; iodine, iodine monochloride, and iodine monobromide; have been studied by Popov and Rygg.l4 This study shows that the acid strength of the halogens toward the pyridine bases decreased in the order 1C1 > IBr > 12. This order agrees with the prediction by Scott15 based on thermo- dynamic considerations that the Lewis acid strength of halogens and interhalogens should follow the order ICl >> BrCl > IBr >> I > Br2> C1 2 2' Steric Considerations Popov and Rygg also14 reported that the stability of the halogen or interhalogen complex with each of the three pyridines decreased in the order 2-picoline > pyri- dine > 2,6-lutidine. The order of base strength of the pyridine to the aqueous proton is 2,6-lutidine > 2-picoline > pyridine. This reversal of ordering was attributed to steric hindrance. 16 studied the interaction Chaudhuri and Basu between iodine and a few methyl substituted pyridines in chloroform. They took the pKa (-£og of the dissociation constant of the conjugate acid) values of the pyridines as being proportional to their first ionization potentials and claimed that the pKa's were related to the formation constants of the iodine complexes although they had too few experimental points to determine the exact nature of this relationship. They found that the 2,6-1utidine- iodine and the 2,4,6—collidine-iodine systems deviated significantly. This deviation was attributed to steric effects. Similar results were reported for the 2,6-lutidine- iodine complex by Bhaskar and Singh.17 Halleux18 reported that in the reaction between phenol and pyridine bases C6H50H + B v_ c6H50H--—B the strength of the hydrogen bonds formed varies in the following order: 3,5-1utidine > 4-picoline > 2,6-lutidine > 2-picoline > pyridine. The basic strength of the com- pounds is, 2,6-lutidine > 3,5-lutidine ~ 4-picoline ~ 2— picoline > pyridine. The above—mentioned investigations showed that steric hindrance is the main factor in the observed dis- crepancies between the basic strength and strength of the complexes. In the determination of basic strength, the very small proton is involved, and there is little steric l9 effect present. Brown and Mihm reported the absence of any important steric effects in the addition of the proton to pyridine, 2,6-lutidine, and monoalkylpyridines. Sacconi, et al.,20 studied the heats of neutralization in water of a series of pyridine bases which include pyridine, picolines and lutidines. They found that for pyridine, picolines, and lutidines a linear relationship between the heats of neutralization and basicity constants is followed, indicating the absence of any steric effect in the reactions. For the large iodine or interhalogen molecules, the presence of substituent groups adjacent to nitrogen atom will invariably hinder the combination of the Lewis acid with the base, and thus weaken the molecular complex formed. Donor Strength The work by Chaudhuri and Basu16 was a step toward evaluating the effect of donor strength on the stability of the complex. Their results were not very meaningful, however, since they evaluated the formation constants of only four pyridine-iodine complexes, three of which in- volved steric factors. Bhaskar and Singh17 extended this work by measuring the formation constants of eleven pyri- dine-iodine complexes. They state that the formation con- stants of the complexes vary "roughly" in the same order as the pKa values of the pyridines. A number of investigators have studied the rela- tionship between the basicity of pyridines and their ability to form various types of complexes. A comprehen- sive study of the relationship between pKf (-£og of the 10 formation constant) of silver complexes of amines with their pKa values by Bruehlman and Verhock21 has shown that plots of pKf XE; pKa fall on two straight lines, one for the pyridines and primary aliphatic amines and one for secondary amines. Recently Cattalini and co-workers22 have shown that the plot of log K XE: pKa for the reaction AuC14" + Py :: AuPyClB + c1" gave three straight lines, one for pyridines without steric hindrance, one for pyridines with one methyl group in the 2-position and a third one for pyridines with methyl groups in the 2 and 6 positions. Solvent Effects While numerous authors have studied the influence of the structure of the donor and/or acceptor molecules on the complex—forming reaction, the effect of the prop- erties of the reaction medium has not been carefully scrutinized. Table 1 gives the reported values of the formation constant of the pyridine-iodine complex in M"1 in five different solvents. The values for this quantity vary from a low of 44 in chloroform to a high of 185 in gfheptane. Obviously some of this variance is due to experimental error as is indicated by the range of values in carbon tetrachloride and in g-heptane. Nevertheless, the low value in chloroform does indicate some effect of the solvent on the formation constant of the complex. 11 Table 1.-—Literature Values for the Formation Constant of the Pyridine-Iodine Complex in Several Solvents Solvent Formation Constant ngexane 127 Bhaskar and Singh17 " 106 Alosi, eE_§1.23 n-Heptane 138 Bist and Person24 " 14o Bhaskar and Singh17 " 185 Reid and Milliken9 interpolated from Figure 4 " 108 Mazzucato, eE_el.25 Cyclohexane 126 Bhaskar and Singh17 " 107:25 Make and Plyler26 " 96 Krishna and Chaudhuri27 " 131 Lake and Thompson28 Carbon tetrachloride lOl Popov and Ryggl4 " " lll Bhaskar and Singh17 " " 183 Sobczyk, eE_§1.29 Chloroform 63 Bhaskar and Singh17 " 44 Chaudhuri and Basul6 Merrifield and Phillips 30 studied the complex be- tween tetracyanoethylene (TCNE) and three benzene deriva- tives. The complexes were examined in three solvents, dichloromethane, diethyl ether, and chloroform. 12 The difference in the equilibrium constants was explained by invoking a solvent competition for the TCNE. Formation constants for the TCNE-solvent interaction were deter- mined relative to an assumed value of zero for the TCNE- chloroform complex. The values for the benzene and sub- stituted benzene complexes were then corrected for this interaction, and the apparent discrepancy between the equilibrium constants was eliminated within experimental error. Solvent competition for one of the adducts was also used by Foster and Hammick31 to explain the results they obtained on the N,N-dimethylaniline-g-trinitroben- zene system. Some studies have also been carried out on com- plexes which are more polar than the n-n type of inter- action in the TCNE complexes. In two of these cases, the Lewis acid used was iodine. The bases in these two studies 32 and triphenylarsine.3 were N,N-dimethylacetamide Dichloromethane, benzene, dioxane, and 3-methysu1folane were used in the first investigation and dichloromethane, carbon tetrachloride and acetonitrile in the second. In both cases, reactions in solvents with high dielectric constants (3-methylsulfolane and acetonitrile) yielded triiodide ion as a reaction product. While the number of solvents studied was rather small, the results seemed to indicate an increase in the strength of the complex with increasing dielectric constant of the solvent. On the 13 surface, these results would seem to be contrary to Briegleb's34 prediction that complex formation constants would have an inverse functional dependence on the dielec- tric constant of the solvent. However, his prediction was based on the formation of a non-polar complex in which the effect of complexation would be to squeeze out solvent molecules in the solvation sphere of the two reactants. This action should result in a weakening of the complex since the uncomplexed reactants would be more highly solvated than the complex. This effect should increase with increasing dielectric constant of the medium. 17 Bhaskar and Singh previously reported that the forma- tion constant of the pyridine-iodine complex decreased with increasing dielectric constant of the reaction medium. These results cannot be regarded as conclusive for several reasons. In the first place, the authors used gfhexane, gfheptane, cyclohexane, carbon tetrachloride and chloro- form as solvents. The range of dielectric constants, therefore, was very narrow with chloroform having the highest dielectric constant of 4.8, and the values for the other solvents hover around 2. Likewise, the differ- ences in the values of the formation constants were very small, and no attempt was made to correct the results obtained in chloroform for the hydrogen-bonding equilib- rium between the solvent and pyridine.35 EXPERIMENTAL Chemicals Pyridine: Fisher "Certified" pyridine was refluxed over granulated barium oxide for 12 hours and fractionally distilled through a 60 cm Vigreaux column. The distillate was then stored in the dark to prevent photo decomposition. One hundred ml portions of this pyridine were then refluxed over BaO for two hours and fractionally distilled through a 20 cm helices packed column and stored over NaOH as they 36 were needed; bp = 115° (lit. bp = 115.58°). Iodine: The purification of iodine has been de- scribed previously.37 Carbon Tetrachloride: The purification of carbon tetrachloride has been described previously;37 bp76o = 76.8° (lit.36 = 76.75°). bp760 3,4-Lutidine: Aldrich Chemical Company (A.C.C.) was refluxed 12 hours over BaO and fractionally distilled through a one meter helices packed column; n 24 = 1.5099 D (lit.38 nD25 2i6-Dichloropyridine: A.C.C. 2,6-dichloropyridine = 1.5099). was recrystallized twice from a mixture of water and 39 ethanol; mp = 87.0 - 87.2° (lit. mp = 87 - 88°). 14 15 All other pyridines were purchased from A.C.C. except for the 4-E-butyl and 3-ethyl which were gifts of the Reilly Tar and Chemical Company. All of these com- pounds were purified by refluxing over BaO for at least two hours and fractionally vacuum distilling through a one— half meter, helices packed column. The compounds and their physical constants are given in Table 2. Dichloromethane: Fisher "Certified" dichloromethane was refluxed for 24 hours over BaO and distilled through a one meter helices packed column; bp760 = 39.9° (lit.36 — O ngeptane: A.C.C. g-heptane was purified by stirring 32 of the alkane with 300 ml portions of concentrated sul- phuric acid until the acid was only slightly colored after two days exposure to the heptane. The alkane was then washed three times with one liter portions of water, dried for several days over anhydrous calcium sulfate, and frac- tionally distilled through an annular teflon spinning band column at a reflux ration of 6:1; bp760 = 98.3 (lit.36 bp760 = 98.427). Chloroform: Malinkrodt, N. F., chloroform was puri- fied by shaking a portion with 50% of its volume of water sever times. It was then stored in the dark over CaSO4 for at least six hours at which time the water content was less than two millimolar as determined by a Karl Fisher titration. The chloroform was then filtered into a glass stoppered 16 Table 2.--Physical Constants of Several Substituted Pyridines Compound ggggiggi Eitifi2t3§55¥3iie Eif' Constant ence 4-Ethylpyridine n34 = 1.4996 n30 = 1.5010 40 2-Flouropyridine n34 = 1.4658 n30 = 1°4678 41 2-Chloropyridine bp = 167.6—168.0° bp = 167-168° 42 2,5-Lutidine n34 = 1.4981 n35 = 1.4982 43 4-Picoline n34 = 1.5026 n35 = 1.5029 44 3-Chloropyridine bp = l48.5-l49.0° bp = 143.5-148° 45 3-Bromopyridine n34 = 1.5686 n30 = 1.5694 46 2-Bromopyridine n34 = 1.5692 n30 = 1.5713 46 3-Picoline n34 = 1.5036 n34 = 1.5043 47 4-t-Butylpyridine n34 = 1.4934 ngs's = 1.4934 48 3,5-Lutidine n34 = 1.5030 n35 = 1.5032 49 bottle. The increase in water content due to this operation was determined to be insignificant. The solvent was then used immediately. The entire procedure from the beginning of purification until the final measurement never took more than 24 hours; bp760 = 61.1° (lit.36 = 61.152). bp76o ngichlorobenzene: Eastman (99+%) gfdichloroben- zene was agitated for 24 hours with concentrated sulfuric acid, washed three times with 50% of its volume of distilled 17 water, dried over CaSO4 for 24 hours, refluxed over BaO for 12 hours, and fractionally distilled through a one-half meter Vigereaux column; bp760 = 181° (1it.36 = l80.48°). bp760 meichlorobenzene: A.C.C. medichlorobenzene was refluxed over BaO for 24 hours and fractionally distilled through a one meter helices packed column; bp760 = 173° (11t.36 = 173.00°). bP760 ngexadecane: A.C.C. thexadecane was agitated with concentrated sulfuric acid until a new portion of acid was only slightly colored after six hours of contact. It was then washed three times with 50% of its volume of distilled water and dried over CaSO4 for 12 hours. The solvent was then fractionally vacuum distilled from BaO through a one-half meter Vigreaux column: fp = 18.2° 50 (lit. fp = 18.165°). l,l,2-Trifluorotrichloroethane: Matheson Coleman and Bell "Spectroquality" l,1,2-trifluorotrichloroethane was refluxed for 12 hours over BaO and fractionally dis- tilled through a one meter helices packed column at a re- flux ratio of 8:1; bp743 = 46.8° (111:.51 bp760 = 47.6°). Benzene and Toluene: Fisher "99% Mole Pure" ben- zene and Eastman "Sulfur-free" toluene were purified in the following manner. One gallon of the solvent was shaken for five hours with two 500 m1 portions of concentrated sulphuric acid. The first portion of acid removed was slightly colored while the second was colorless. 18 The solvent was then shaken with one liter of water, then with one liter of dilute aqueous potassium hydroxide solution and finally with two, one liter portions of water. The prod- uct was then dried over CaSO4 for 12 hours, and fractionally distilled through a one meter helices packed column. The boiling points of the benzene and toluene are respectively; 36 = 80.0° and llO.3° (lit. = 80.103° and b9750 bp760 110.623°). g-Xylene: A.C.C. p—xylene was purified by the method of fractional freezing. Water content was deter- mined to be less than 3 millimolar in the final product; 37 fp = 13.2° (lit. fp = 13.263°). 1,2-Dichloroethane: 1,2-dichloroethane was puri- 36 fied by the method of Vogel;53 = 83.4° (lit. bp760 bp760 = 83.483°). Nitromethane: Nitromethane was purified by the 36 method of Clarke and Sandler;54 = lOl.2° (lit. bp760 _ 0 Preparation and Stability of Solutions For the study of the complexes between fourteen substituted pyridines and iodine in carbon tetrachloride, solutions of iodine in carbon tetrachloride were prepared by the standard techniques. Although their concentration could be determined by iodometric titrations, it was found that equal accuracy was obtained by measuring the absorption 19 of the solution at 517 mu and calculating the concentra-. tion from the value of the molar absorptivity at this wave- length (6 = 927). Stock solutions of the respective 517 pyridines were made by weighing the respective compound into a volumetric flask and diluting with purified solvent. Mixed solutions of iodine and of the respective pyridine were prepared just before each measurement. Con- tact of solutions with the atmosphere was kept to a minimum but no attempt was made to do all the transfers in a com- pletely inert atmosphere. Preliminary results have indicated that brief contacts of the solution with the atmosphere did not alter the experimental data. Likewise it was found that solutions of iodine and the pyridines were stable for at least one hour after mixing and, in general, spectral measurements were completed within 5-10 minutes of the initial mixing. For the study of the pyridine-iodine complex in different solvents, stock solutions of pyridine and iodine were prepared by weighing a quantity of the reactant into a clean, dry volumetric flask. They were then diluted to volume with the appropriate solvent in a water-bath ther- mostated at 25.0°. The prOper aliquots of these solutions were then pipetted into another volumetric flask which was brought to volume in a thermostated bath kept at the same 20 temperature at which the final absorption measurements were made. Solutions of iodine and pyridine were prepared just prior to each measurement. Again contact of solutions with the atmosphere was kept to a minimum, but no attempt was made to do all of the transfers in an inert atmosphere. The results indicate that the brief contacts of the solu- tions with air did not alter significantly the experimental data. It was found that pyridine slowly catalyzed the dehydrochlorination of chloroform and 1,2-dichloroethane. In these cases, special care was taken to use solutions as soon as possible after their preparation. A study of absorbance ye. time for a pyridine-iodine solution in the above solvents showed that the decomposition was incon- sequential during the time required to complete the mea— surements. Solutions of pyridine and iodine in nitro— methane were also found to be very unstable. Spectral Measurements In the study of the substituted pyridine-iodine complexes, absorption spectra were obtained on a Cary Model 14 Spectrophotometer. Either 1 cm or 5 cm path- length cells were used. All measurements were carried out at room temperature of ~ 25°. The donor/acceptor ratio was varied within such limits as to give a good spread of experimental points. Spectra of iodine-amine 21 solutions containing a fixed amount of iodine and a variable amount of the amine were obtained in the 530-390 mu spectral range. A typical set of absorption curves is illustrated in Figure 1. All pyridine derivatives listed in Table 29 gave essentially ideal isobestic points. In the case of 4-phenylpyridine and of 4-cyan0pyridine, a marked shift in the isobestic point was observed with a change in the amine/iodine ratio. It seems quite likely that this behavior may be due to the formation of weak Py.212 complexes since second molecule of iodine may form a weak bond with the phenyl or the cyano group, respec- tively. In the study of the pyridine-iodine complex in different solvents spectral measurements were made on a Beckman DU spectrophotometer with a thermostated cell com- partment. The wavelength scale of the instrument was calibrated using a holmium oxide filter. The accuracy of the absorbance scale was checked by the method of additive absorbances and by the method of G. HauptS4 using standard alkaline chromate solutions. The temperature of the cell compartment was determined by inserting a calibrated .thermometer into the compartment through a styrofoam cover and allowing the system to come to equilibrium. The temperature could be controlled to i 0.l°. 22 Figure 1.--Absorption spectra of the pyridine—iodine system in carbon tetrachloride solutions; CI = 4.44 x 2 10’4 M; ch (M): (1) 0.0, (2) 0.00381, (3) 0.00762, (4) 0.00834, (5) 0.01143, (6) 0.01524, (7) 0.01669, (8) 0.02286, (9) 0.02503; (10) 0.03338. 23 o 0. ”024010004. «0 «o 400 500 520 M?" It) 420 24 Calculations Spectral data obtained on the donor-iodine mix- tures (Tables 3-27) were used to calculate the formation constant of the complex using the method described by Ketelaar, et al.55 In Ketellar's equation 1 1 l l =-——- +———- (1) e -e C K Ts —e ) e -e t 12 Py M c 12 c I2 8t is the apparent molar absorptivity of 12 (i.e., the measured absorbance of the solution divided by the total 12 concentration), 6c and BI are the molar absorptivities 2 of the complex and the I respectively, K is the formation 2 M constant of the complex in 9 mole-1 and CPy is the total concentration of base moles/t. A plot of l/(et - £1 ) ye. 2 l/C should give a straight line. From the lepe and FY intercept of the line the formation constant and the molar absorptivity of the complex can be determined. It should be noted that this method is likely to give better results than the treatment of Benesi and Hildebrand3 since mea- surements had to be made at wavelengths at which both the complex and iodine absorb. A regression analysis of the data was performed on a CDC 3600 computer with points over three standard devia- tions off the line being rejected (Appendix I). Large excesses of pyridines were found to cause a shift in the 25 isobestic point due to an increase in the polarity of the solvent which favors the formation of the triiodide ion. Thus, it was necessary in most cases to use base concentra- tions such that it was no longer possible to equate the concentration of the free base with the total concentra- tion of the base which is a necessary condition for the use of Ketellar's equation. To overcome this difficulty an iteration was performed on the equation in the follow- ing manner. The experimental values of l/CPy and l/(st - 612) were used in Ketelaar's equation to calculate a value for the concentration of the complex which was then subtracted from the value of CPy to give a new array. The process was then repeated until successive values for KM agreed within 0.1%. In performing this type of mathematical analysis a great deal of caution must be exercised since relatively small errors in certain experimental parameters can cause large errors in the evaluation of the formation constant. A good illustration of such an error is the extreme sen- sitivity of the obtained K value to the iodine concentra- M tion for complexes where KM < 10. In this study it was found that an error of 2.0% in the iodine concentration caused a 50% change in the value for KM of the 2-flouro- pyridine complex. This uncertainty is evident in the large standard deviation for the K of the complex. M 26 The iteration procedure for solving Ketelaar's equation offers a distinct advantage in the evaluation of formation constants from experimental data. The values converged rapidly and reproducibly at all wavelengths. This method makes it possible to use lower concentrations of the donors and thus it avoids the possibility of a change in the physical properties of the solvent in solu- tions with high concentrations of the base. .It also makes it unnecessary to use a complex iterative procedure such as the one reported by Conrow, e£_el.56 As in all calculations of this type, we assumed that activity coefficients are equal to unity, i.e., the values of K- M should be noted that this assumption may be rather danger- are expressed in concentration units. It ous especially in the case of weak complexes where a large excess of donor (or acceptor) has to be added in order to produce a measurable change in the absorption spectrum.57 The concentration of the component in excess may be as high as 2 - 4 M, and under these conditions it would be unreasonable to expect that the activity coefficients, even of uncharged species, can be ignored. Calculations were carried out at least at three different wavelengths over at least a 20 mp range. It has been previously shown that the linearity of the Benesi- Hildebrand plot by itself is not a sufficient criterion for assuming a simple 1:1 interaction between the donor and the acceptor molecules.58 27 In the study of solvent effects on the formation constants of the pyridine-iodine complex the values of the 1 formation constants were changed from M- to mole fraction units by use of the following relationship: KM Kx = v +x (v -v ) (2) S Py Py S where KX and KM are the mole fraction and the molar equilib- rium constants, X is the mole fraction of pyridine and FY VPy and VS are the molar volumes of pyridine and of the solvents. The above equation is based on the assumption that no volume change occurs on mixing. In the present study the concentration of pyridine was always kept below 0.03 M and the concentration of iodine was ~ 10-4 M. Therefore, X < 0.01 and under these conditions, equation PY (1) can be reduced to KX = KMS° (3) where S° is the concentration of pure solvent in moles/2. Whenever possible, corrections were made for the solvent-solute interaction. It has been shown that chloro- form forms a hydrogen—bonded complex with pyridine3S with Kx = 0.69. Likewise, it is well known that aromatic com- pounds form charge-transfer complexes with iodine. In fact, the formation constants of benzene, toluene, and 28 xylene complexes with iodine have been reported to be 0.15, 0.16, and 0.31 I?..mole"l respectively.59 In order to allow for the solvent-solute interac- tion in the case of the four solvents mentioned above, the corrected values of the formation constants were calculated from the expression32’60 (see Appendix 2) Kcorr = Kobs (l + KS S ) (4) where KS is the equilibrium constant (M-l) for the solute- solvent interaction, K is the experimental value for obs pyridine-iodine complex in the given solvent and 8° is the concentration of the pure solvent in moles/R. Enthalpy and entropy values for the pyridine-iodine complex were calculated in the usual manner by determining the slope and intercept of a plot of in K XE- l/T by a M least squares procedure. Measurements of the absorption maximum of the blue- shifted iodine band of the complex were made in a series of mixed solvents obtained by adding varying amounts of nitromethane to carbon tetrachloride. The dielectric con- stants of these mixed solvents were calculated by use of the pr1nc1ple of add1t1v1ty (i.e., Dcalc = XCC14 DCC14 + XCH3NOZDCH3N02)' The value of dielectric constant of nitromethane used in the aforementioned calculations was 35.9. The results of these measurements are contained in Figure 5. Table 3.--Experimental Data on the 2—F1uoropyridine-Iodine Complex in Carbon Tetrachloride pr A450mu A440nm A430mu ~ 25°, [12] 1.64 . 10‘4 M, path length 5.000 cm 0.1220 0.291 0.240 0.195 0.0697 0.225 0.172 0.130 0.1917 0.359 0.312 0.263 0.1568 0.325 0.276 0.229 0.0666 0.225 0.170 0.129 0.0686 0.224 0.171 0.129 0.0961 0.276 0.223 0.178 0.1163 0.284 0.234 0.189 0.1647 0.332 0.282 0.234 0.2085 0.369 0.322 0.271 0.1099 0.250 0.199 0.153 0.0880 0.260 0.205 0.161 812 150.0 82.0 44.0 KM 2.33.8 2.03.7 1.73.8 e 1070 1110 1130 C 30 Table 4.--Experimental Data on the 2,5-Lutidine-Iodine Complex in Carbon Tetrachloride pr A420nm A410nm A400mu ~ 25°, [12] = 6.50 - 10-5 M, path length = 5.000 cm 0.00398 0.290 0.294 0.268 0.00597 0.342 0.350 0.328 0.00796 0.374 0.381 0.347 0.00995 0.399 0.405 0.368 0.01194 0.417 0.421 0.383 0.01593 0.440 0.446 0.407 0.00546 0.331 0.336 0.305 0.00818 0.379 0.384 0.350 0.01091 0.409 0.414 0.371 0.01364 0.428 0.433 0.393 0.01637 0.440 0.446 0.407 0.02182 0.459 0.464 0.422 0.00651 0.350 0.358 0.329 0.00976 0.397 0.404 0.370 0.01301 0.420 0.431 0.398 0.01626 0.439 0.450 0.412 0.01951 0.453 0.463 0.424 0.02602 0.468 0.479 0.440 812 20.9 10.2 6.41 KM 30112.0 30412.2 30817.4 6 1630 1660 1510 C 31 Table 5.--Experimenta1 Data on the 3—Chloropyridine-Iodine Complex in Carbon Tetrachloride pr A440mu A430mu A420mu ~ 25°, [12] = 2.19 - 10'4 M, path length = 5.000 cm 0.03290 0.575 0.577 0.546 0.04935 0.706 0.721 0.689 0.06580 0.800 0.828 0.792 0.08225 0.873 0.910 0.876 0.09870 0.926 0.966 0.930 0.02775 0.520 0.517 0.484 0.04163 0.645 0.654 0.620 0.05551 0.736 0.753 0.719 0.06939 0.809 0.839 0.806 0.08326 0.871 0.905 0.867 0.02126 0.444 0.431 0.401 0.03180 0.560 0.561 0.530 0.04240 0.640 0.647 0.614 0.05301 0.720 0.736 0.702 0.06361 0.779 0.801 0.768 0.03439 0.578 0.580 0.549 e1 82.0 44.0 20.9 KMZ 17.21.25 16.51.28 16.41.29 6 1290 1400 1360 C Table 6.-—Experimental Data on the 2-Chloropyridine-Iodine Complex in Carbon Tetrachloride CPy A440mu A430mu A420mu ~ 25°, [12] = 2.20 - 10-4 M, path length = 5.000 cm 0.03011 0.224 0.170 0.124 0.04517 0.277 0.220 0.168 0.06022 0.330 0.270 0.209 0.09033 0.427 0.351 0.280 0.02132 0.187 0.137 0.096 0.04265 0.268 0.214 0.160 0.06397 0.340 0.281 0.218 0.08530 0.401 0.340 0.268 0.12790 0.513 0.448 0.359 0.03578 0.243 0.189 0.142 0.05366 0.308 0.249 0.190 0.07155 0.364 0.302 0.239 0.08944 0.418 0.350 0.279 0.03272 0.231 0.177 0.131 0.06543 0.349 0.289 0.226 s 82.0 44.0 20.9 I2 KM 3.81.15 3.51.13 3.11.10 cc 1240 1200 1110 33 Table 7.~-Experimental Data on the 2-Brom0pyridine-Iodine Complex in Carbon Tetrachloride m CPy A430mu A420mu A410mu A400mu ~ 25°, [12] = 2.23 - 10"4 M, path length = 5.000 cm 0.03274 0.192 0.144 0.100 0.068 0.04912 0.250 0.191 0.137 0.091 0.06549 0.298 0.230 0.167 0.110 0.08186 0.348 0.272 0.199 0.131 0.09823 0.391 0.303 0.226 0.151 0.02804 0.171 0.126 0.090 0.059 0.04206 0.221 0.166 0.118 0.079 0.05608 0.271 0.210 0.150 0.099 0.07010 0.313 0.242 0.173 0.116 0.08412 0.351 0.274 0.199 0.131 0.03062 0.185 0.137 0.094 0.062 0.06124 0.290 0.223 0.161 0.107 0.07655 0.333 0.261 0.189 0.126 0.09186 0.373 0.292 0.217 0.143 0.02465 0.160 0.113 0.080 0.053 0.04931 0.251 0.190 0.138 0.091 0.09861 0.396 0.312 0.230 0.155 51 44.0 20.9 10.2 6.41 KMZ 4.511.8 4.31.26 4.41.21 4.41.24 e 1040 870 640 420 C 34 Table 8.--Experimental Data on the 4-Ethylpyridine-Iodine Complex in Carbon Tetrachloride pr A420mu A410nm A400mu ~ 25°, [12] = 6.50 -10"5 M, path length = 5.000 cm 0.03280 0.465 0.487 0.460 0.00570 0.302 0.313 0.293 0.01139 0.390 0.403 0.379 0.00285 0.217 0.224 0.211 0.05110 0.480 0.503 0.473 0.00888 0.359 0.374 0.350 0.01776 0.421 0.439 0.414 0.00222 0.189 0.194 0.181 0.00444 0.273 0.282 0.264 0.00588 0.305 0.318 0.299 0.01178 0.384 0.402 0.375 0.02355 0.443 0.463 0.437 0.02944 0.455 0.477 0.449 0.04122 0.470 0.491 0.462 0.00364 0.248 0.257 0.240 0.00729 0.336 0.349 0.327 0.01459 0.406 0.424 0.398 0.08390 0.491 0.517 0.491 e 20.9 10.2 6.41 I2 KM 25111.9 24811.5 24611.7 9 1590 1660 1570 C 35 Table 9.--Experimenta1 Data on the 3-Bromopyridine-Iodine Complex in Carbon Tetrachloride pr A435mu A430nm A425mu A420m ~ 25°, [12] = 1.096 - 10‘4 M, path length = 5.000 cm 0.02286 0.249 0.248 0.191 0.231 0.03428 0.316 0.314 0.309 0.299 0.04571 0.370 0.371 0.367 0.356 0.05714 0.409 0.412 0.409 0.397 0.06857 0.443 0.448 0.444 0.431 0.09142 0.491 0.500 0.497 0.483 0.02058 0.232 0.230 0.173 0.214 0.03086 0.294 0.293 0.289 0.280 0.04115 0.347 0.347 0.342 0.331 0.05144 0.388 0.390 0.386 0.373 0.06172 0.420 0.423 0.420 0.409 0.08230 0.473 0.479 0.475 0.462 0.01615 0.198 0.193 0.188 0.179 0.02422 0.254 0.252 0.247 0.237 0.03230 0.301 0.301 0.296 0.286 0.04037 0.341 0.342 0.339 0.327 0.04845 0.377 0.379 0.373 0.361 0.06460 0.427 0.432 0.428 0.416 21 62.3 44.0 30.2 20.9 KM2 17.51.19 17.41.22 17.11.18 17.21.15 e 1430 1460 1470 1430 C 36 Table lO.--Experimental Data on the Pyridine-Iodine Complex in Carbon Tetrachloride pr A430mu A420mu A415mu A400mu ~ 25°, [12] = 6.50 - 10'5 M, path length = 5.000 cm 0.01209 0.267 0.288 0.290 0.06952 0.403 0.441 0.447 0.01046 0.248 0.267 0.268 0.01395 0.279 0.301 0.303 0.01744 0.300 0.327 0.329 0.02093 0.319 0.345 0.349 0.02790 0.347 0.376 0.378 0,01295 0.270 0.292 0.295 0.01943 0.308 0.335 0.339 0.02590 0.338 0.367 0.371 0.03238 0.355 0.390 0.394 0.03886 0.368 0.403 0.407 0.05181 0.387 0.423 0.427 0.17000 0.425 0.470 0.477 0.07392 0.401 0.443 0.449 0.02957 0.349 0.382 0.386 eI 44.0 20.9 14.4 KM2 11111.3 10911.1 10811.0 cc 1380 1530 1550 ~ 25°, [12] = 4.44 - 10-4 M, path length = 1.000 cm 0.05154 0.579 0.585 0.518 0.02577 0.501 0.508 0.445 0.07732 0.608 0.616 0.541 0.10310 0.622 0.631 0.558 0.15460 0.645 0.653 0.582 0.01016 0.357 0.359 0.312 0.02032 0.469 0.473 0.412 0.03048 0.521 0.527 0.462 0.04064 0.552 0.558 0.489 0.05080 0.574 0.581 0.510 0.06096 0.588 0.594 0.521 0.00381 0.200 0.202 0.179 0.00762 0.287 0.288 0.254 0.01143 0.373 0.380 0.337 0.01524 0.402 0.407 0.360 0.01905 0.459 0.465 0.409 Table 10 . --Continued 37 C Py A430mu 420mu A415mu A400mu 0. 02286 0.466 0.470 0.414 0- 00834 0.309 0.311 0.276 0- 01669 0.426 0.430 0.380 0. 02503 0.487 0.489 0.431 0. 03338 0.503 0.512 0.451 51 25.0 13.8 6.25 2 KM 10612.3 10812.6 11112.7 cc 1510 1520 1330 Table ll.--Experimental Data on the 4-E-Butylpyridine- Iodine Complex in Carbon Tetrachloride CPy A420mu A4 15mm A410mu A405mu A400mu ~ 25°, [12] = 5.86 - 10‘4 M, path length = 5.000 cm 0.000732 0.157 0.160 0.159 0.156 0.149 0.001464 0.258 0.223 0.276 0.271 0.259 0.002195 0.351 0.361 0.363 0.358 0.342 0.002927 0.416 0.428 0.432 0.427 0.409 0.00 3659 0.470 0.485 0.489 0.483 0.462 0.004391 0.513 0.531 0.537 0.529 0.505 0.005854 0.579 0.600 0.607 0.599 0.572 0.000947 0.189 0.192 0.192 0.189 0.181 0.001894 0.312 0.323 0.326 0.320 0.306 0.002842 0.406 0.417 0.421 0.413 0.397 0.003789 0.473 0.488 0.492 0.486 0.464 0.004736 0.527 0.542 0.548 0.540 0.517 0.005683 0.569 0.589 0.593 0.584 0.560 0.007578 0.632 0.656 0.662 0.653 0.625 0.001166 0.219 0.222 0.224 0.220 0.209 0.002331 0.355 0.366 0.368 0.362 0.346 0.003497 0.453 0.469 0.473 0.465 0.444 0.004662 0.517 0.536 0.541 0.532 0.510 0.005828 0.568 0.587 0.595 0.585 0.560 0.006993 0.611 0.632 0.639 0.630 0.602 0.009324 0.670 0.695 0.703 0.692 0.662 51 20.9 14.4 10.2 7.19 6.41 K142 30114.6 29614.7 29614.6 30014.1 29914.0 5 1560 1630 1650 1620 1550 38 Table 12.--Experimenta1 Data on the 4-Picoline-Iodine Complex in Carbon Tetrachloride CPy A420mu A415mu A410mu A405nm ~ 25°, [12] = 4.50 - 10-4 M, path length = 1.000 cm 0.00614 0.419 0.429 0.430 0.421 0.01229 0.529 0.545 0.547 0.534 0.01843 0.587 0.602 0.607 0.594 0.02457 0.614 0.635 0.639 0.627 0.03072 0.638 0.657 0.661 0.650 0.00289 0.278 0.283 0.283 0.276 0.00579 0.399 0.407 0.408 0.399 0.00869 0.470 0.481 0.482 0.471 0.01159 0.512 0.527 0.528 0.517 0.01448 0.546 0.560 0.561 0.549 0.01738 0.569 0.587 0.589 0.576 0.00405 0.328 0.336 0.337 0.328 0.00811 0.451 0.462 0.463 0.452 0.01217 0.517 0.531 0.532 0.520 0.01622 0.552 0.569 0.571 0.558 0.00711 0.431 0.445 0.446 0.433 0.01423 0.541 0.556 0.558 0.548 eI 20.9 14.4 10.2 7.19 KMZ 21913.7 21613.4 21613.5 21513.8 e 1600 1650 1660 1630 C 39 Table l3.--Experimental Data on the 3-Picoline-Iodine Complex in Carbon Tetrachloride CPy A420ml A415mu A410mu A405mu ~ 25°, [12] = 5.83 - 10-4 M, path length = 1.000 cm 0.00378 0.411 0.419 0.418 0.407 0.00755 0.572 0.585 0.584 0.568 0.01134 0.647 0.671 0.671 0.655 0.01512 0.709 0.727 0.728 0.708 0.01890 0.748 0.766 0.766 0.748 0.02268 0.771 0.791 0.791 0.771 0.00292 0.345 0.351 0.350 0.339 0.00584 0.502 0.512 0.512 0.498 0.00877 0.598 0.608 0.608 0.591 0.01169 0.659 0.673 0.673 0.657 0.01462 0.692 0.709 0.710 0.689 0.01754 0.731 0.747 0.747 0.728 0.00203 0.268 0.270 0.269 0.261 0.00407 0.412 0.420 0.419 0.408 0.00610 0.511 0.520 0.517 0.500 0.00814 0.570 0.582 0.580 0.561 0.01018 0.618 0.631 0.630 0.611 0.01221 0.657 0.670 0.668 0.649 0.00461 0.438 0.447 0.445 0.430 0.00923 0.600 0.610 0.609 0.591 eI 20.9 14.4 10.2 7.19 KMZ 20712.8 20612.9 20713.0 20713.4 e 1600 1640 1640 1590 C 40 Table 14.--Experimental Data on the 3,5-Lutidine-Iodine Complex in Carbon Tetrachloride pr A415mu A410ml A405mu A400mu ~ 25°, [12] = 5.81 - 10'4 M, path length = 1.000 cm 0.00279 0.505 0.511 0.505 0.483 0.00559 0.675 0.685 0.679 0.651 0.00838 0.773 0.788 0.782 0.754 0.01118 0.823 0.837 0.831 0.803 0.01397 0.862 0.886 0.880 0.854 0.01677 0.899 0.920 0.919 0.894 0.02235 0.942 0.969 0.972 0.952 0.00081 0.222 0.223 0.221 0.211 0.00163 0.363 0.368 0.362 0.349 0.00102 0.264 0.267 0.262 0.251 0.00204 0.421 0.426 0.421 0.402 0.00306 0.533 0.542 0.536 0.515 0.00408 0.606 0.615 0.609 0.586 0.00510 0.661 0.672 0.665 0.640 0.00612 0.707 0.717 0.711 0.685 0.00816 0.767 0.779 0.772 0.743 0.00293 0.517 0.523 0.519 0.499 0.00586 0.690 0.701 0.697 0.670 0.00880 0.785 0.799 0.792 0.767 s1 14.4 10.2 7.19 6.41 KM2 38212.3 37812.8 37713.3 36913.6 e 1760 1800 1790 1740 C 41 Table 15.--Experimental Data on the 3,4‘LUtidine-Iodine CPy Complex in Carbon Tetrachloride A415ml A410nm A405mu A400mu ~ 25°, [12] = 6.54 - 10"4 M, path length = 1.000 cm 0.00223 0.00335 0.00446 0.00558 0.00669 0.00186 0.00279 0.00372 0.00465 0.00558 0.00162 0.00243 0.00323 0.00404 0.00485 0.00253 0.00379 0.00506 6 K12 M e C 0.517 0.637 0.716 0.763 0.815 0.458 0.576 0.654 0.717 0.771 0.413 0.525 0.605 0.673 0.717 0.530 0.645 0.728 14.4 427i10.2 1700 0.526 0.648 0.726 0.779 0.831 0.466 0.585 0.665 0.728 0.787 0.416 0.532 0.614 0.681 0.730 0.536 0.655 0.740 10.2 4Zlill.4 1740 0.523 0.645 0.724 0.774 0.827 0.464 0.582 0.663 0.726 0.783 0.412 0.528 0.609 0.675 0.727 0.532 0.652 0.734 7.19 420i12.3 1730 0.504 0.624 0.701 0.744 0.800 0.446 0.560 0.641 0.702 0.754 0.395 0.509 0.591 0.653 0.797 0.511 0.625 0.709 6.41 386:17.9 1740 42 Table l6.--Experimental Data on the 3-Ethy1pyridine—Iodine Complex in Carbon Tetrachloride pr A420mu A415mu A410mu A405mu ~ 25°, [12] = 6.59 - 10'4 M, path length = 1.000 cm 0.00908 0.711 0.731 0.732 0.716 0.01816 0.849 0.872 0.875 0.854 0.00208 0.337 0.342 0.342 0.334 0.00415 0.505 0.521 0.522 0.507 0.00623 0.617 0.631 0.633 0.617 0.00831 0.688 0.707 0.708 0.689 0.01038 0.739 0.758 0.759 0.738 0.01246 0.740 0.759 0.760 0.740 0.01661 0.831 0.855 0.857 0.835 0.00449 0.528 0.538 0.538 0.526 0.00898 0.704 0.723 0.723 0.706 0.01348 0.792 0.812 0.814 0.795 0.01797 0.842 0.866 0.869 0.848 0.02246 0.878 0.902 0.905 0.881 0.00118 0.221 0.224 0.223 0.217 0.00236 0.364 0.370 0.370 0.360 0.00354 0.462 0.474 0.474 0.461 0.00471 0.537 0.549 0.549 0.532 0.00707 0.647 0.660 0.661 0.644 0.00943 0.712 0.731 0.731 0.715 61 20.9 14.4 10.2 7.19 KM2 24311.0 24310.9 24511.0 24711.4 e 1580 1620 1620 1580 C 43 Table l7.--Experimental Data on the Pyridine-Iodine Complex in Carbon Tetrachloride pr A410mu A416mu A420mu A428nm 8.5°, [12] = 4.362 - 10‘4 M, path length = 1.000 cm 0.00528 0.369 0.371 0.368 0.346 0.00793 0.440 0.443 0.438 0.408 0.01057 0.486 0.487 0.483 0.448 0.01322 0.515 0.519 0.511 0.472 0.01586 0.543 0.549 0.539 0.498 0.02115 0.575 0.581 0.569 0.527 0.00619 0.398 0.402 0.398 0.370 0.00928 0.466 0.470 0.464 0.432 0.01238 0.511 0.515 0.507 0.468 0.01548 0.539 0.542 0.534 0.492 0.01857 0.566 0.569 0.561 0.517 51 9.2 16.0 18.3 36.7 KM2 21611.5 21311.9 22011.9 22612.1 EC 1610 1630 1600 1450 25.0°, [12] = 5.383 ‘4 M, path length = 1.000 cm 0.01522 0.498 0.496 0.480 0.03043 0.621 0.616 0.597 0.04565 0.675 0.671 0.648 0.06087 0.702 0.697 0.672 0.07609 0.717 0.691 0.01447 0.488 0.486 0.470 0.02171 0.564 0.560 0.539 0.02895 0.611 0.606 0.586 0.00724 0.347 0.348 0.336 0.03618 0.643 0.638 0.613 0.04342 0.662 0.662 0.636 e1 11.1 18.6 30.0 KMZ 10410.4 10510.2 10310.5 6 1510 1500 1450 C Table 17.--Continued 44 CPy A410mu A416ml A420mu A428ml 16.3°, [12] = 4.362 - 10‘4 M, path length = 1.000 cm 0.00663 0.342 0.346 0.346 0.00996 0.409 0.414 0.413 0.01326 0.459 0.463 0.458 0.01989 0.512 0.515 0.511 0.02652 0.553 0.555 0.546 0.00874 0.389 0.395 0.391 0.01312 0.453 0.457 0.452 0.01749 0.503 0.509 0.503 0.02186 0.528 0.532 0.527 0.01658 0.490 0.493 0.489 0.02623 0.549 0.553 0.546 51 11.5 16.0 18.3 KMZ 15111.5 15311.8 15911.8 cc 1580 1580 1550 34.5°, [12] = 5.084 - 10‘4 M, path length = 1.000 cm 0.01854 0.430 0.442 0.442 0.426 0.02781 0.503 0.519 0.519 0.495 0.03708 0.553 0.566 0.565 0.538 0.04635 0.587 0.589 0.598 0.571 0.05562 0.610 0.621 0.620 0.593 0.01414 0.379 0.391 0.392 0.378 0.02827 0.503 0.520 0.520 0.496 0.04241 0.573 0.388 0.588 0.563 0.02121 0.457 0.466 0.467 0.450 0.03534 0.545 0.556 0.557 0.531 0.05655 0.611 0.620 0.620 0.594 81 11.8 17.7 23.6 39.3 KM2 68.810.48 70.810.48 70.810.46 71.410.48 e 1510 1530 1530 1450 C Table l7.—-Continued 45 CPy A410mu A416mu A420m.) A428mu 25.0°. [12] = 5.383 - 10‘4 M, path length = 1.000 cm 0.00617 0.317 0.327 0.327 0.310 0.01234 0.452 0.462 0.459 0.433 0.01851 0.529 0.541 0.538 0.512 0.02468 0.583 0.595 0.591 0.563 0.03085 0.613 0.624 0.619 0.589 0.01767 0.523 0.533 0.530 0.504 0.00532 0.279 0.296 0.296 0.286 0.01065 0.425 0.435 0.433 0.404 0.01597 0.501 0.514 0.511 0.485 0.02130 0.562 0.571 0.569 0.539 0.02662 0.596 0.606 0.603 0.571 0.00776 0.360 0.369 0.368 0.354 0.01551 0.500 0.513 0.511 0.485 0.02327 0.574 0.588 0.585 0.556 51 13.0 20.4 22.3 42.7 KMZ 10011.2 10710.6 10810.7 10511.0 e 1520 1520 1500 1430 C 46 Table 18.--Experimental Data on the Pyridine-Iodine Complex in Dichloromethane CPy A390mu A395mu A400mu A410mm 25.0°, [12] = 5.390 - 10"4 M, path length = 1.000 cm 0.001265 0.150 0.155 0.155 0.149 0.002530 0.257 0.266 0.267 0.249 0.003796 0.343 0.352 0.352 0.325 0.005061 0.407 0.420 0.419 0.383 0.006326 0.460 0.472 0.472 0.432 0.004505 0.382 0.393 0.392 0.364 0.006070 0.452 0.463 0.462 0.430 0.007509 0.507 0.520 0.519 0.479 0.009011 0.550 0.565 0.562 0.517 0.003325 0.320 0.327 0.327 0.302 0.004988 0.412 0.422 0.422 0.388 0.006650 0.476 0.487 0.486 0.451 0.008313 0.530 0.544 0.542 0.499 CI 9.3 9.3 11.1 27.8 KMZ 14910.8 15310.5 15210.6 15110.6 cc 1800 1820 1820 1670 16.4°, [12] = 3.640 - 10‘4 M, path length = 1.000 cm 0.00494 0.344 0.349 0.343 0.307 0.00741 0.411 0.415 0.408 0.364 0.00989 0.455 0.459 0.449 0.401 0.01236 0.488 0.492 0.482 0.428 0.01484 0.514 0.518 0.506 0.447 0.01978 0.553 0.555 0.542 0.485 0.00640 0.382 0.386 0.382 0.348 0.00960 0.445 0.450 0.442 0.399 0.01281 0.487 0.492 0.485 0.437 0.01921 0.537 0.543 0.532 0.476 eI 13.7 13.7 16.5 33.0 KMZ 21314.1 21713.9 22213.0 22313.6 e 1850 1860 1810 1610 C Table 18.--Continued 47 CPy A390ml A395ml A400mu A410mu 8.7°, [12] = 4.059 - 10‘4 M, path length = 1.000 cm 0.00394 0.426 0.431 0.425 0.00591 0.497 0.505 0.496 0.00789 0.550 0.557 0.546 0.00986 0.579 0.587 0.576 0.01184 0.608 0.613 0.601 0.00318 0.384 0.388 0.380 0.00635 0.514 0.519 0.507 0.00953 0.580 0.584 0.571 0.00476 0.464 0.469 0.463 0.00794 0.550 0.558 0.548 0.01272 0.618 0.620 0.612 e1 7.4 9.9 14.8 KM2 33912.8 34011.8 33913.5 cc 1880 1890 1860 32.6°, [12] = 5.100 10‘4 M, path length = 1.000 cm 0.00503 0.310 0.322 0.323 0.301 0.00754 0.400 0.411 0.411 0.384 0.01006 0.464 0.476 0.475 0.444 0.01258 0.520 0.530 0.530 0.491 0.01510 0.562 0.575 0.572 0.530 0.02013 0.629 0.641 0.638 0.591 0.00601 0.333 0.345 0.346 0.325 0.00901 0.436 0.447 0.447 0.418 0.01202 0.502 0.516 0.516 0.476 0.01202 0.503 0.517 0.517 0.477 0.01803 0.580 0.594 0.595 0.545 eI 13.7 15.7 21.6 33.3 KM2 9612.4 10112.4 10212.3 10312.3 8 1850 1850 1840 1690 C 48 Table l9.--Experimental Data on the Pyridine-Iodine Complex in 1,2-Dichloroethane CPy A385nm A390mu A395mu A404mu 25.2°, [12] = 4.585 - 10‘4 M, path length = 1.000 cm 0.00370 0.351 0.358 0.358 0.339 0.00555 0.430 0.441 0.441 0.419 0.00740 0.498 0.511 0.511 0.480 0.00925 0.545 0.559 0.559 0.525 0.01110 0.581 0.596 0.596 0.561 0.01481 0.646 0.657 0.656 0.606 0.00430 0.387 0.394 0.394 0.368 0.00645 0.471 0.480 0.480 0.448 0.00860 0.535 0.548 0.547 0.509 0.01076 0.582 0.596 0.595 0.552 0.01291 0.616 0.628 0.626 0.584 61 10.9 13.1 15.3 26.2 KMZ 19013.0 18712.3 18712.2 19211.2 cc 1890 1940 1940 1790 pr A390nm A395nm A400mu A410nm 8.4°, [12] = 3.624 - 10-4 M, path length = 1.000 cm 0.00358 0.401 0.406 0.394 0.352 0.00537 0.464 0.469 0.456 0.407 0.00716 0.507 0.512 0.500 0.443 0.00895 0.536 0.540 0.529 0.466 0.01075 0.559 0.562 0.550 0.483 0.00295 0.367 0.372 0.365 0.326 0.00590 0.483 0.487 0.477 0.424 0.00885 0.537 0.542 0.529 0.467 0.00443 0.439 0.441 0.431 0.383 0.00738 0.515 0.518 0.506 0.447 0.01181 0.570 0.575 0.562 0.492 e1 13.8 16.6 22.1 44.2 KMZ 40913.5 41612.8 41214.8 42413.6 1 1900 1900 1860 1630 C Table 19.--Continued 49 pr A390mu A395mu A400mu A410mu l6.3°, [12] = 3.851 - 10‘4 M, path length = 1.000 cm 0.00491 0.405 0.410 0.407 0.366 0.00736 0.471 0.479 0.472 0.427 0.00982 0.525 0.529 0.522 0.467 0.01473 0.577 0.584 0.571 0.515 0.00608 0.429 0.437 0.431 0.390 0.01217 0.523 0.531 0.525 0.470 0.01826 0.586 0.594 0.585 0.525 0.00913 0.510 0.517 0.507 0.449 0.01522 0.579 0.587 0.574 0.508 0.02435 0.628 0.635 0.618 0.546 eI 13.0 15.6 18.2 41.5 KM2 266112.6 268111.4 280111.l 285111.0 cc 1860 1880 1830 1620 33.3°, [12] = 4.440 - 10-4 M, path length = 1.000 cm 0.00547 0.357 0.357 0.349 0.321 0.00821 0.438 0.439 0.429 0.392 0.01095 0.505 0.508 0.496 0.449 0.01369 0.548 0.550 0.538 0.485 0.01643 0.581 0.585 0.571 0.517 0.00626 0.387 0.387 0.379 0.345 0.01253 0.528 0.529 0.519 0.468 0.01879 0.607 0.607 0.595 0.536 0.00939 0.474 0.474 0.461 0.420 0.01253 0.532 0.531 0.522 0.470 0.02506 0.667 0.666 0.648 0.583 61 11.3 13.5 18.0 38.3 K 2 13411.6 13311.4 13211.2 13411.0 cc 1920 1930 1890 1680 50 Table 20.—-Experimental Data on the Pyridine-Iodine Complex in Chloroform pr A414nm A404mu A400mu A395mu 25.1°, [12] = 5.355 - 10‘4 M, path length = 1.000 cm 0.00504 0.238 0.251 0.249 0.242 0.00756 0.313 0.326 0.326 0.322 0.01009 0.368 0.393 0.393 0.382 0.01261 0.416 0.411 0.441 0.429 0.01513 0.451 0.482 0.482 0.469 0.02017 0.512 0.548 0.547 0.530 0.00814 0.328 0.349 0.348 0.337 0.01221 0.412 0.439 0.438 0.426 0.01628 0.461 0.492 0.490 0.480 0.02035 0.511 0.548 0.546 0.531 0.02442 0.545 0.583 0.582 0.569 e1 20.5 11.2 9.3 9.3 KM2 79.010.6 77.710.7 77.010.7 77.010.5 cc 1550 1670 1680 1640 33.3°, [12] = 5.730 - 10‘4 M, path length = 1.000 cm 0.00760 0.271 0.282 0.280 0.271 0.01140 0.354 0.369 0.366 0.356 0.01520 0.414 0.435 0.431 0.419 0.01900 0.461 0.486 0.483 0.468 0.02280 0.500 0.528 0.524 0.508 0.03041 0.565 0.595 0.592 0.570 0.01078 0.339 0.355 0.354 0.341 0.01616 0.427 0.449 0.448 0.430 0.02155 0.488 0.516 0.512 0.496 0.02694 0.538 0.568 0.565 0.546 0.03233 0.577 0.607 0.603 0.586 eI 19.2 8.7 8.7 7.0 KM2 57.010.3 56.610.1 56.610.2 56.810.2 e 1540 1640 1630 1580 C Table 20.--Continued 51 CPy A414mu A404mu A400mu A395mu 16.4°, [12] = 5.064 - 10‘4 M, path length = 1.000 cm 0.00566 0.313 0.335 0.341 0.332 0.00849 0.394 0.427 0.430 0.421 0.01132 0.447 0.484 0.487 0.479 0.01415 0.488 0.529 0.531 0.525 0.01698 0.523 0.570 0.571 0.563 0.02265 0.569 0.616 0.622 0.615 0.00748 0.365 0.394 0.400 0.394 0.01123 0.447 0.481 0.488 0.481 0.01497 0.487 0.539 0.545 0.536 0.01872 0.531 0.586 0.591 0.582 0.02246 0.560 0.615 0.620 0.612 61 21.7 11.8 9.9 7.9 KM2 12211.6 11611.0 12110.7 11710.8 to 1520 1690 1680 1680 8.7°, [12] = 4.551 - 10‘4 M, path length = 1.000 cm 0.00698 0.375 0.414 0.420 0.416 0.01048 0.442 0.489 0.496 0.491 0.01397 0.477 0.537 0.543 0.541 0.01746 0.510 0.575 0.581 0.580 0.02095 0.533 0.602 0.609 0.607 0.02794 0.567 0.638 0.645 0.640 0.00923 0.413 0.463 0.469 0.465 0.01384 0.477 0.537 0.543 0.538 0.01864 0.521 0.580 0.587 0.579 0.02307 0.551 0.610 0.618 0.611 0.02769 0.569 0.632 0.640 0.633 e1 19.8 11.0 8.8 6.6 KM2 17413.6 16911.5 17211.6 17112.3 6 1500 1690 1710 1690 C 52 Table 21.--Experimenta1 Data on the Pyridine-Iodine Complex in M-Heptane pr A430nm A424mu A420mu A410mp 26.2°, [12] = 4.699 - 10‘4 M, path length = 1.000 cm 0.00563 0.301 0.304 0.301 0.275 0.00844 0.366 0.371 0.368 0.336 0.01126 0.408 0.415 0.410 0.377 0.01409 0.443 0.450 0.447 0.412 0.01689 0.469 0.478 0.473 0.437 0.02252 0.503 0.511 0.509 0.467 0.00617 0.315 0.320 0.316 0.291 0.00925 0.383 0.386 0.385 0.353 0.01234 0.426 0.427 0.425 0.395 0.01543 0.455 0.463 0.461 0.429 0.01851 0.481 0.488 0.487 0.452 e1 29.8 17.0 12.8 6.4 KM2 15211.0 15511.0 15311.2 14911.7 cc 1380 1400 1400 1300 32.3°, [12) = 4.699 ’4 M, path length = 1.000 cm 0.00916 0.335 0.337 0.333 0.303 0.01374 0.398 0.395 0.397 0.362 0.01832 0.440 0.446 0.440 0.404 0.02290 0.469 0.476 0.471 0.431 0.02784 0.489 0.498 0.494 0.453 0.03664 0.525 0.532 0.528 0.489 0.01059 0.357 0.360 0.357 0.330 0.01589 0.419 0.424 0.420 0.388 0.02118 0.463 0.469 0.465 0.427 0.02648 0.490 0.497 0.492 0.452 0.03177 0.509 0.519 0.513 0.466 eI 40.4 29.8 23.4 17.0 KMZ 11411.0 11011.6 11210.8 11211.8 e 1380 1410 1390 1280 C Table 21.--Continued 53 pr A430mu A424nm A420mu A410mu l6.5°, [12] = 4.002 - 10‘4 M, path length = 1.000 cm 0.00980 0.392 0.404 0.401 0.373 0.01470 0.434 0.445 0.444 0.414 0.01960 0.464 0.477 0.476 0.444 0.02450 0.480 0.492 0.491 0.460 0.02940 0.487 0.504 0.504 0.468 0.03920 0.511 0.526 0.526 0.488 0.01027 0.398 0.408 0.408 0.380 0.01540 0.438 0.450 0.449 0.422 0.02053 0.466 0.480 0.479 0.446 0.02566 0.486 0.500 0.500 0.465 0.03080 0.492 0.508 0.507 0.472 :1 40.0 27.5 22.5 15.0 KMZ 23114.3 23114.2 22814.1 23113.5 so 1410 1450 1450 1350 7.9°, [12] = 4.002 - 10‘4 M, path length = 1.000 cm 0.00528 0.372 0.384 0.384 0.361 0.00792 0.422 0.434 0.434 0.408 0.01057 0.454 0.466 0.466 0.442 0.01321 0.470 0.487 0.487 0.460 0.01585 0.483 0.498 0.498 0.471 0.02113 0.502 0.520 0.520 0.488 0.00674 0.403 0.418 0.418 0.392 0.01011 0.446 0.458 0.458 0.434 0.01348 0.475 0.488 0.488 0.458 0.01685 0.490 0.508 0.509 0.477 0.02022 0.501 0.518 0.518 0.485 61 40.0 27.5 22.5 17.5 KM2 36213.8 36614.8 46815.2 36914.7 6 1420 1470 1470 1380 C 54 Table 22.-~Experimental Data on the Pyridine—Iodine Complex in 1,l,2-Trifluorotrichloroethane CPy A420mu A414ml A410mu 25.0°, [12] = 5.566 - 10‘4 M, path length = l.0000conz Aommv pmuomaom 0H03 mmoam> ummm .mme .muumH80cu pomHmmm one whom mo cows: HmcomumcuoucH .mcowuoaom mooonwm cw momwm oecmmuo mo mucmumcoo coquHOOmmeom .nouaom Nae memmo.o-mo~oo.o om.m o.anee~ am.e maaoanmaasaumum woe aoooo.o-~meoo.o oe.o o.~ana~e em.o oeaoanmaasnuosaoue.m woe eamao.oueamooo.o ma.o m.mneam Hm.m oaaoanaoamnumsaoum.m oae oammo.oueomoo.o aa.m e.anmem omo.o oaaoanaoaanumue Mae moomo.oumamoo.o oe.o ~.~heom omo.o maeoanaoaanuosaoum.~ oae emmaoo.oueoeaoo.o aa.m m.ehoom om.m oaaoanmnaaasmuoue flee Neomo.ouaa~oo.o ma.m m.mnea~ om.e oaaoanaaaanquIe Nae oaaao.oumamoo.o mo.m o.mnao~ ma.m oaaoanaoaanhozum mae ooee.o-moeo.o am.m m.anaoa omo.o oaaoanae ewe meaao.o-maoao.o ea.~ o~.ohm.ea oa.a manoenaaosonmum one oaaao.o-o~a~o.o em.~ om.ono.oa aa.~ oaaoanaoonoanoum n Homao.o-mee~o.o oa.o mm.one.e m~.~ oneoanaaosonmum n aema.o-~mamo.o me.o ma.one.m o~.~ oaaoanmaonoanoum n mmo~.o-eamo.o ee.o- m.oho.~ eo.a moaoanmaonosaaum a va aa.e oaaoanaoonoenoaoIo.~ x08 .ocou 000m 0M0 ucmumcoo 2 OH x NH mo mmcom wmwpflod M wo .0000 when 0cm mcepeumm Suez mmxmameoo mcwpoH mo 6mm n 00 mcoHuoaom opwuoacomuume conumu ca mocwpeuam pmuoueumnsm UHOEQV mucmumcoo 00a002H0m11.m~ canoe 63 .uco>HOm once «0 cOADMHucmocoo on» mmEHu muwcs “waoe CH m mOHOE 0H mm mo 00H0> one .muwcs cowuomum 0HoE 0e cm>em me x 0» H0000 aaoumsexoummm we mg no moaow 0:90 .Amoaao pea .eo ..Eono .6111 .e .neenn .3 .2 one ..no .aoamxnom .e .eo .Ammmav oomv .vn ..Oom .EOSU .Ed .0 .HOMOOM .2 .m can m3OHUG¢ .6 .AQ ..o.m~ on negate Heat aae mo.m H.10NH1 H.~haam manomonxomue ooe ma.a mahaemm a.~haom maonemnonoenoaouo eoe eo.m mmheoaa a.~hm- moonaononoenoaous mam om.oa amhammm N.thme maneuoonoanoaoum.a mam ao.a m.anaem~ o.oaama mannuosonoanoao eae ae.m m.ahm~ma a.onama homeomonoanoennonosaeaneum.H.a oae e~.~ ~.en~moa e.onmoa moenoenonnuou nonnnw ewe Na.a H.10eooa ~.Hnama manuommua ooe am.e knee o.oao.o e.enmoa o.oaa.ae anemonoaao eoe m~.~ omem nma.o o.mamma m.oha.~m oaonamm moe mm.~ mmom noa.o FHHNNa m.anm.em mamoaoe Noe e~.~ 11mm nam.o m.eaame a.oho.ea memaaxtm xnsa o unooxs me xx was naoeaom mucm>dom DGOHOMMHQ ca xmamaoo ocepoHlmcwpeumm may no mucmumcou cowumeuomll.om manna 64 There is a considerable amount of disagreement between our data and those of Bhaskar and Singh.17 In general, their values of the formation constants are about 50% smaller than the values listed in Table 1. It should be noted that Bhaskar and Singh carried out their measurements in chloroform solutions. Chloroform, be- sides being more polar than carbon tetrachloride, is known to hydrogen bond to pyridine.35 The specific ef- fects of these two factors are discussed in the following section. Chloroform is also a notoriously poor solvent for the study of halogen complexes due to the ease with which it is oxidized by air to give HCl as one of the products. Hydrochloric acid easily reacts with halogens to give polyhalide ions, and this reaction, obviously, can introduce a large error in the spectrophotometric study of iodine complexes. Chloroform can only be used as a solvent for such studies if it is very carefully purified just prior to use. It also appears that the spectral measurements of Bhaskar and Singh17 were car- ried out only at a single wavelength. It has been shown by Person and co-workers that it is possible to correlate the formation constants of halogen complexes with acetonitrile and chlorinated 61 where 0* is a acetonitrile with the Taft 0* constant, measure of only the inductive effect in a rigid aliphatic compound. Since in the case of substituted pyridines the 65 resonance effects would likewise be important, it seemed more reasonable to investigate the possible correlation between the Hammett 0 function and the log of the forma- tion constant of the respective complexes. A plot of log KM X§° 0 is shown in Figure 2. It is seen that a reasonably straight line is obtained. The equation of the line is given by log KM = —2.250 + 2.11 The standard deviation of the lepe is 0.16 and the value 62 is 0.98. of r, Jaffe's correlation coefficient, According to Jaffe's criteria62 for the magnitude of the correlation coefficient, the value of r indicates good agreement with the Hammett equation. A similar study of iodine complexes of substituted styrylpyridines by Mazzucato, eg_el.,63 gave values of p = -2.1 with r = 0.93 (values for pyridine, 4-methylpyridine and 3—methylpyridine were also included in the calculation). Rather poor cor- relation as shown by the low value of r may be due to the interaction of the halogen with the delocalized electrons of the stilbene ring (c.f. our results with the 4-phenyl- pyridine and 4-cyanopyridine described above). There are some rather minor discrepancies which illustrate the limitations of the theory. For example, one would expect the value of 0 for 3-chloropyridine to be greater than that of 3-bromopyridine which would imply that 66 Figure 2.--Relationship between log KM and Hammett 0 function for the pyridine-iodine complexes in carbon tetra- chloride solutions. 67 l 0.. m._ oar: mo. Wm 68 the KM value for the corresponding 3-bromopyridine complex would be the larger of the two. Experimentally the values for K follow the eXpected order, while the 0 values do not. M It was also of interest to us to determine if there is a correlation between the acid dissociation constants of the pyridines and the iodine complex formation constant. As seen from Figure 3, a plot of log K ye. pKa does indeed M yield a fairly straight line although there is some scatter of the experimental points. This is implied from the com- bination of our results in Figure 3 and those of Jaffé and Doak64 in which they obtained a linear plot between the pKa's of several substituted pyridines and the Hammett 0 constant. The advantage of plotting the log of the for— mation constant of the complex ye. pKa for the pyridines is that the values for sterically hindered pyridines can be compared. From this comparison one can gain an under- standing of the relative importance of steric factors. As can be seen in Figure 3 there is very little difference between the steric hindrance for the proton and the iodine molecule for pyridines substituted in the two positions. The effect of steric factors on the 2,6-lutidine-iodine complex is quite pronounced. It can be concluded, therefore, that in pyridine and in substituted pyridines, in the absence of steric effects, there is a definite parallelism between the com- plexing abilities of the amines and their basicities. 69 Figure 3.-—Re1ationship between the acidity constant of the pyridines in aqueous solutions and the formation constants of the pyridine-iodine complexes in carbon tetrachloride solutions. 70 0.0 0.0 Qt Bu 0 q q 1 u 4 l-N fiv-N. O LQ-N 1°.— . In 0 to-» 02-0.4. o r. e(-« 0 NW...“ 01-0." 0 . 011.0 0 0 0:700 SM 71 Such correlation, of course, would be expected if one ad0pts Lewis' definition of acids and bases since both H+ and I are Lewis acids. It should be noted, however, 2 that such generalizations are only applicable to cases where the reference bases do not differ appreciably in structure. Solvent Effects Formation constants of the pyridine-iodine complex in twelve different solvents are presented in Table 30. Comparison of these values with various physical proper- ties of the respective solvents indicates a possible cor- relation only with the dielectric constant of the medium. A plot of D ye. KX is shown in Figure 4. While the points do not fall on a smooth curve, there seems to be little doubt that, other factors being equal, an increase in the dielectric constant of the solvent results in an increase in the stability of the complex. The observed scatter, in all likelihood, is the result of specific solvent-solute interactions. It seems reasonable to expect that the increase in the bulk dielectric constant of the reaction medium will result in greater stability of the pyridine-iodine complex. The dipole moments of iodine, pyridine, and of the complex 65 are 0.0, 2.20, and 4.90 Debyes respectively. It appears, therefore, that with increasing polarity, the solvents will 72 Figure 4.——Relationship between the formation constant expressed in mole fraction units of the pyridine-iodine complex and the dielectric constant of the solvent in which it was measured. x 73 x 000m OOON 000. a q _ o. e. 1 I: 001000 nzonroo 1000 0 . 0 0 e on m. an 6 .oo - :18 o-.. «noonuoo r o 1.. noro 0 0 1 «oermoé News . «stereo-.. n N 0 .o .66 to 0 74 tend to solvate more strongly the highly polar complex and thus contribute to its stability. Kobinata and Nagakura,66 as well as Boule,67 recently reported a study of the dipole moments of iodine complexes with aliphatic amines in ben— zene and in dioxane solutions. All of the values obtained were in the 6.0 - 7.2 Debye range. Kobinata and Nagakura noted that the dipole moments of the complexes increased with increasing concentration of the amines in the solvent mixtures. Since the dielectric constant of the medium likewise increased with the amine concentration, the .authors interpreted the results by postulating an increase in the stabilization of the charge-transfer. These re- sults agree with our observation that in solvents with D > 15, the charge separation in the complex increases to the extent that clear cut separation of the charges occurs with the resulting formation of the triiodide ion. The above phenomena would be expected from an examination of the theoretical behavior of polarizable dipoles in a changing dielectric medium.68 Indeed, other ramifications of this behavior would be expected. Accord- ing to Mulliken's theory9 of charge-transfer complexes, the spectrum of the complex would be altered by the dielectric constant of the medium. In particular, the energy of the charge-transfer band would be shifted. The direction of this shift would depend upon the relative importance of two factors, if one assumes that the energy 75 of the no bond wave function remains constant. One effect would be due to the increase in energy of the dative wave function as a result of the increased charge separation, and the other effect would be due to the decrease in energy of the dative wave function as a result of the solvation of the polar species. Due to the absorption of pyridine and of the solvents in the ultraviolet region, the charge-transfer band cannot be observed. Lacking this information it is nevertheless informative to examine the relationship between the dielectric constant of the solvent and the wavelength of maximum absorption of the blue shifted iodine band of the complex as shown in Figure 5. The blue shifted iodine band has been explained by Mulliken.4 This explanation is based on the assumption that in the excited state of the complex an electron has been transferred from the donor to the iodine molecule. This electron goes into an antibonding orbital of the iodine which is diffuse and increases the effective size of the iodine molecule. When light is absorbed by the complexed iodine molecule, the excitation energy is sup- plemented by an energy of repulsion between the donor and the abnormally large iodine molecule. The blue shift should become larger with increasingly close contact be- tween the donor and iodine which is related to the bond strength of the complex. Thus the fact that the frequency shift of the iodine band is related to the dielectric 76 Figure 5.--Relationship between the wavelength of maximum absorption of the blue-shifted iodine band of the complex and the dielectric constant of the solvent in which it was measured. Solvents: (l) e-heptane; (2) e—hexadecane; (3) carbon tetrachloride; (4) XCCl4 = 0.99284 and XCHBNOZ = 0.00716, Dcalc = 2.48; (5) 1,1,2-trifluorotrichloroethane; (6) XCCl = 0.9857 4 and XCH3N02 = 0.0143, Dcalc = 2.72; (7) XCCl4 = 0.9646 and XCH3N02 = 0.0345, Dcalc = 3.43; (8) XCC14.= 0.9304 and XCH3N02 = 0.0696, Dcalc = 4.58; (9) chloroform; (10) dichloromethane; (11) 1,2-dichloroethane; (12) toluene; (l3) benzene; (l4) E-xylene; (15) M-dichloro— benzene; (l6) e—dichlorobenzene. 77 .0} .N O.N 0.0 o 0.0. 78 constant of the medium is another indication of the influence of the dielectric constant on the complexation reaction. It is also interesting to note the deviation of the points (12— 16) in cases where aromatic solvents were used. This can easily be understood in terms of the known specific solute- solvent interactions which have already been mentioned. Thermodynamic data obtained by measuring the forma- tion constant of the complex as a functiOn of temperature are given in Table 31. As can be seen by examining the values, the differences are not significant within experi- mental error and, therefore, do not, at this time, contribute to our understanding of the role of the solvent in the com- plexation reaction. On the basis of the above results it seems reasonable to conclude that the reports on the decrease in the stability of iodine complexes with increasing dielectric constant of 17 cannot be correct. the medium It is clear, however, that it would be naive to expect a monotonic correlation between a given physical property of a series of solvents and the stability of the pyridine-iodine complex since this approach neglects spe- cific solvent-solute interactions. While we endeavored to correct in the case of aromatic solvents and chloroform, this correction is tenuous at best, and it only takes into account the interaction of iodine with the solvent in the first three cases and of pyridine with the solvent in the 79 Table 31.--Enthalpy and Entropy Values for the Pyridine-Iodine Complexation in Different Solvents ===r Solvent -AH(kcal/mole) -AS(e.u.) n-Heptane 8.16:0.22 17.3i0.8 Carbon tetrachloride 7.47i0.06 15.8i0.2 Dichloromethane 8.59:0.31 18.9:l.l 1,2-Dichloroethane 7.77:0.16 15.6i0.6 Chloroform 7.82:0.19 l7.5:0.6 last. There is little doubt that weak interactions exist between other solvents and pyridine as well as iodine and that these interactions influence overall stability of the complex. In fact it has been shown recently that there is a significant interaction between pyridine and carbon tetrachloride (as well as other aromatic compounds).69 In the calculation of formation constants, however, it is'a common practice to arbitrarily treat the solvent as an inert dispersing medium regardless of the magnitude of the concentrations of the solutes. In certain cases it may be possible to find systems where such assumptions may be justifiable,70 but such cases must be rather exceptional. In general, it seems safe to conclude that if we wish to understand the role of the solvent in molecular complex formation, we should know the nature and the extent of solute-solvent interactions with both reactants and products to the maximum accuracy attainable. See Adden- dum (p. 104). PART II ALKALI METAL IONS IN PYRIDINE AND ACETONE SOLUTIONS HI STORI CAL INTRODUCTI ON Solvation The importance of ion-solvent interactions in solutions of electrolytes cannot be overemphasized. Yet at this time these interactions are understood only in rather crude qualitative terms. Recent deve10pments in the use of far-infrared spectroscopy in the study of solu- tions of alkali metal salts (ammonium salts are considered to be part of this group) in several nonaqueous solvents have added a new perspective from which these interactions can be examined. Several authors have reported the appearance of new bands in this region of the spectrum which have been described primarily in terms of cation- solvent vibrations although some anion effects on these bands have been noted. The work presented in this portion of the thesis is an extension of the aforementioned studies in two other solvents, pyridine and acetone. A more extensive historical discussion of solvation studies can be found in the Ph.D. theses of Brian W. Maxey71 and John L. Wuepper.72 81 82 Pyridine The donor properties of pyridine toward a wide variety of Lewis acids have been well established; however, its usefulness as a solvent has not been studied extensively. Pyridine is characterized by a wide liquid range (-4l.8° t0~ 115.6°), a low dielectric constant of 12.3, a moderate dipole moment of 2.20 Debyes, and a Trouton constant of 21.8, indicating that it is relatively unassociated in the liquid phase. The donor prOperties of pyridine and its moderate dipole moment should make pyridine a good solvating agent for alkali metal ions although the solu- bility of these salts may not be particularly high due to the low dielectric constant of the solvent. Two solid complexes between pyridine and lithium chloride have been found by Brussed and Halut-Desportes in their study of the solubility of lithium chloride in 73 pyridine. They found that a solid complex corresponding to the formula LiCl-CSHSN existed in contact with a satu- rated solution of the salt in pyridine above a temperature of 19.3°. Below this temperature the solid corresponded to the trisolvate, LiCl°3C5H5N. The latter compound melts incongruently at 19.3 i 0.l°. Slow evaporation of a solu- tion of LiCl in pyridine open to the laboratory atmosphere produced a solid with a composition of LiCl°H20-2CSHSN. The existence of interactions between alkai metal ions and pyridine was used by Burgess and Kraus74 as an explanation 83 for the low conductances of these ions in pyridine. This study also gave the ion pair dissociation constants for lithium, sodium, potassium, and ammonium picrates as 4 4 4 0.83 x 10' , 0.43 x 10' , 1.00 x 10'4, and 2.8 x 10‘ , respectively, and the ion pair dissociation constants for sodium, potassium, and ammonium iodides as 3.7 x 10-4, 4, and 2.4 x 10-4, respectively. The values of 2.1 x 10' these constants indicate a high degree of ion pairing as would be expected for a solvent with a low dielectric constant. The ion pair dissociation constant was found to be relatively unchanged upon addition of water to the pyridine. Good agreement is found between the results obtained by Burgess and Kraus and the more recent conduc- tance work of Mandel and co-workers.75 The infrared and Raman spectra of pyridine have been studies by numerous workers. From infrared and Raman 76 data Corrsin, et al., made a complete assignment of all the pyridine fundamentals. This assignment has been slightly modified by McCullough, gt_al.,77 so that the values of thermodynamic functions statistically calculated from the observed vibrational frequencies would agree with experimental values. The change in assignment made by McCullough has since been criticized and modified by 78 on the basis of an infrared Wilmshurst and Bernstein and Raman study of pyridine and three partially deuterated pyridines. 84 The effect of the pyridine interaction with hydrogen bonding solvents on the infrared spectrum of pyridine has been studied by Takahaski and co-workers.79 They found that there were relatively large shifts to higher energies for several of the skeletal vibrational bands of the pyri— dine molecule. If there were no changes in the electron distribution in the pyridine molecule upon hydrogen bonding, the skeltal vibrations would shift to lower frequencies due to the increased mass. The increase in frequency of these vibrations indicates a considerable change in the electron distribution which strengthen the chemical bonds of the ring system. Upon comparing the spectra of the hydrogen bonded pyridine molecule to the spectra of the pyridinium ion, the authors conclude that the electron distribution of the hydrogen bonded pyridine is close to that of the pyridinium ion. French and Wood80 recently reported a study of ammonium, d4-ammonium, sodium, and potassium tetraphenyl- borates in pyridine solutions by far-infrared techniques. The sodium salt was also examined in three other solvents, 1,4-dioxane, piperidine, and tetrahydrofuran. From the data obtained in this study, the authors claim that sodium tetraphenylborate exists in pyridine solutions as an un- solvated ion pair. This hypothesis seems to be contrary to previous work done in this laboratory and a further investigation of pyridine solutions of alkali metal salts was thought to be worthwhile. 85 Acetone Acetone had been used extensively as a solvent for organic materials. It also has the ability to solvate some inorganic salts as is indicated by the high solu- bility of 0.427 mole fraction units of lithium per- chlorate.81 The solvent is characterized by a wide liquid range (-95.4 to 56.2°), a moderate dielectric constant of 20.70, a relatively high dipole moment of 2.72 Debyes, and a Trouton constant of 21.5 indicating that it is a relatively unassociated liquid. Solubili— ties of electrolytes are enhanced by the dipole moment and the low Trouton constant of acetone. The solid diacetonate of lithium bromide has been prepared by Bell, et al.82 It was found to decompose into the unsolvated salt at 35.5°. On the other hand, the authors found that lithium chloride did not form a solid solvate. The acetonates of sodium iodide and bi- sulfite are well known due to their use in a method for 36 In this method a solu- the purification of acetone. tion of one of the sodium salts in acetone is cooled causing the acetonate to precipitate. The precipitate is then filtered from the mother liquor. Upon heating the precipitate melts and the acetone is fractionally distilled from the solution. A conductance study of lithium bromide solutions in acetone by Olson and Konecny83 gave a value of 86 2.56x10‘4 for the ion pair dissociation constant for the salt. The ion pair dissociation constant for potassium iodide in acetone was determined to be approximately 9x10“3 by Dippy.84 These dissociation constants illustrate that, in acetone, as would be expected for a solvent with a dielectric constant of 20, dissolved salts of alkali metals exist primarily as ion pairs. The infrared spectra of sodium iodide,85 sodium perchlorate,86 and lithium perchlorate81 solutions in acetone have been studied previously from 4000 to 400 cm-1. The primary emphasis in these studies has been on the changes in the infrared spectrum of acetone upon addition of the above salts. The results obtained by these authors are interpreted in terms of complex formation between the cation and the carbonyl group of acetone. The shifts in the acetone bands have also been found to be independent of the anion present. The appearance of a band in the 1 420-430 cm- region of lithium perchlorate-acetone solu- 81 and is attri- 1 tions is mentioned by Pullin and Pollock buted to a higher frequency component of the 380 cm— acetone band. Yamada86 has also studied the effect of lithium and sodium perchlorate on the n+n* transition of acetone in the ultraviolet spectral region, and found that the absorption band for this transition is shifted to higher energy upon addition of the two salts. The effect is more 87 pronounced in the case of the lithium prechlorate. These results are interpreted in terms of charge transfer com- plex formation between the alkali metal ion and the ace- tone. Assuming this type of interaction, the author calculates the percent covalent character of the oxygen- lithium and oxygen-sodium bonds using Pauling's "bond 7 length" relationship.8 The percent covalent character of the bonds is found to be 13% and 8% respectively. EXPERIMENTAL Chemicals Pyridine: Fisher "certified" pyridine was frac- tionally distilled from granulated barium oxide through a one meter, helicies packed column. It was then stored over nolecular sieves (Fisher type 4A). The water content of the pyridine was determined to be approximately three millimolar by a Karl Fisher titration. Acetone: Baker, N. F., acetone was dried over calcium sulphate. The water content as determined by a Karl Fisher titration was approximately seven millimolar. dG—Acetone: Diaprep, Inc., d6-acetone with a minimum isotopic purity of 99.5 percent was dried over molecular sieves (Fisher type 4A) and decanted. Piperidine: Fisher "certified" piperidine was refluxed over granulated barium oxide for two hours and fractionally distilled. Tetrahydrofuran: Fisher "certified" tetrahydro- furan was dried over calcium sulphate and decanted. The water content of the solvent was approximately one milli- molar as determined by a Karl Fisher titration. 1,4-Dioxane: Fisher "certified" 1,4-dioxane was dried over calcium sulphate and decanted. The water 88 89 content of the solvent was approximately one millimolar as determined by a Karl Fisher titration. Alkali Metal Salts: All of the alkali metal salts, with the exceptions of LiI, KBPh NH BPh ND I, and the 4’ 4 4' 4 Li6 salts, were reagent grade chemicals and were used after drying without further purification. Lithium iodide, 98% pure, was obtained from K 8 K Laboratories and was used without purification. The tetra- phenylborates were prepared by adding a stoichiometric amount of sodium tetraphenylborate to an aqueous chloride solution of the desired cation. The precipitate was then washed with water and vacuum dried. The d4-ammonium iodide was purchased as the 98% isotopically pure salt from Dia- prep, Inc., and was used without further purification. The separated isotope, Li6, was purchased as the metal from Union Carbide Oak Ridge Laboratory, Oak Ridge, Tennessee. The assay furnished with the metal showed that it was 6 and 4.4% Li7. In order to prepare the L16 salts, 95.6% Li the metal was first added to water. The resulting basic solution was then neutralized with the reagent grade acid of the desired anion. The titration was followed potentio» metrically. The water was then removed from the solution of the L16 salt by evaporation. All the salts except the Li6I were then dried at 230°. The LiGI was dissolved in acetone and precipitated as the acetonate by cooling in dry ice slush. The acetonate was then decomposed to the 90 pure salt by placing it in a vacuum oven 80° and 0.1 mm pressure for two days. Preparation of Solutions Both of the solvents and most of the salts used in this study were hygroscopic and care was taken to prepare them in as nearly anhydrous conditions as possible. Ex- posure of the solvents and solutions to the atmosphere during preparation or transfer was minimized by performing these Operations with syringes. All solutions were pre- pared at room temperature of approximately 22°. Instrumentation Infrared spectra were obtained on two instruments, a Perkin-Elmer Model 225 Spectrometer with a range of 1 4000-200 cm“ and Perkin-Elmer Model 301 Spectrometer with a range of 666-14 cm-1. Normally the Model 225 instrument 1 and the Model was used for spectra from 4000 to 600 cm- 301 was used in the 666 to 50 cm.1 region. In cases where duplicate measurements were made, they always agreed within experimental error. The frequency scale of the 301 spectrometer was calibrated using the rotational spectrum of water vapor. The construction, optical layout, and electronics of both spectrophotometers are described very well in the Operational manual supplied by the manufacturer. 91 Some suggestions for special instrument procedures on the 301 are given by B. W. Maxey.71 Experimental Techniques Spectra on the 225 Spectrometer were obtained using conventional solution cells with sodium chloride windows in the 4000-600 cm.1 region. Polyethylene spacers were used to obtain pathlengths of 0.015 mm to 0.1 mm. For spectra below 600 cm-1 polyethylene cells with 0.10 mm pathlengths purchased from Barnes Engineering were used. The 225 Spectrometer was operated in the double beam mode with air as a reference. Spectra on the Model 301 were obtained using polyethylene cells purchased from Barnes Engineering with pathlengths of 0.10 or 0.20 mm. The 301 Spectrometer was usually operated in the double beam mode with an equal thickness of solvent in the reference beam. In cases in which this procedure was inconvenient or impossible, an empty polyethylene cell was placed in the reference beam and corrections were made for solvent absorption. Concentrations of salts in the solvents varied between 0.05 and 1.5 M. RESULTS AND DI SCUSSION Infrared Spectra of Alkali Metal Salt Squtions in Pyridine The infrared spectra of alkali metal salt solutions in pyridine have been examined for three effects: 1. shifts in skeletal vibrations of pyridine 2. appearance of a solvent-cation band 3. splitting of perchlorate bands and the appear- ance of mean active bands of this anion due to symmetry lowering by solution structure An examination of the effect of the alkali metal salts on three skeletal vibrations of pyridine yielded results which were completely analogous to the work of Takahashi, §£_§l.,79 mentioned previously. The results are given in Table 32. The skeletal vibrations of pyridine l at 1581, 991.5 and 603 cm- all shift to higher energies upon addition of the alkali salts. The magnitude of this shift is Na+> pr.12 and CS >> CPyS where the superscript zero denotes the analytical concentration of the specie, we can write the following equilibrium ex- pressions. K - CPY'IZ (3) _ _ o _ corr (ng CPy-12)(EPy CPyS) C nyS K = _ 5 (4) S (ng pr.s) CS An expression for the total absorbance of a solution at a particular wavelength, A can also be written, assuming t! a 1 cm pathlength. 120 121 A = C C + €I2 (Ci2 - CPy°12) (5) Inherent in the above equation is the assumption that the pyridine-solvent complex and the pyridine do not absorb light at the wavelength at which the measurements have been made. Substituting in equations (3), (4), and (5) so as to eliminate the concentration of the two complexes we arrive at equation (6). O 1 - 1 + KSCS + 1 (6) € - C I K C5 (C . - e ) e . - e t 12 corr Py Py 12 12 Py 12 12 This equation is identical to the Ketelaar Equation except for the term 1 + KSCS’ which is a constant for low con- centrations of pyridine and iodine. Thus, if we equate 1/K from the lepe of the Ketelaar plot with the term obs in equation (6) O l/Kobs = T— ‘7’ COI‘I‘ and rearrange we obtain = o Kcorr (1 + KSCS)Kobs (8) where K is the formation constant of the pyridine- corr iodine complex corrected for the competing equilibrium. For Case 2, the situation in which the solvent interacts with the specie in limiting concentration, in 122 this instance the iodine, we can follow an entirely analogous develOpment. The equilibria involved are given as follows. Py + 12.1—.Py'I2 (9) 12 + S 1__ 12's (10) ° ‘ O 0 Making the assumptions that CPy >> CPy-IZ and CS >> C12.S we can write the equilibrium expressions given below. C P1012 K = fi— (11) corr (C° - C . - C . YSC° 12 Py 12 12 S Py CIZ'S K = o _ _ o (12) S (C12 CPy°I2 CIZ'SYCS A similar equation for the total absorbance of a solution which includes the previous assumptions can be written as before. A = E C + 6 (CE - C - C (13) . . ) 2 2 2 FY I2 I2 5 Eliminating the concentration of the two complexes in equations (11), (12), and (13) we arrive at an equation which is identical to equation (6). Thus corrections for solvent competition can be made in the same manner as in Case 1. APPENDIX 3 SUGGESTIONS FOR FUTURE WORK 1. Studies of the infrared spectra of polynuclear anions, e.g., CIO-, SCN-, and N03, should be extended in both acetone and pyridine solutions. These studies might give some indication of the degree of anion participation in the solvation shell of the alkali metal ions. 2. An examination of the far-infrared spectra Of solutions of alkali metal salts in several substituted pyridines should give some idea of the effect of donor strength and steric effects on the solvation of the alkali metal ions. 3. The ultraviolet spectrum of solutions of alkali metal salts in pyridine should be examined. The shifts of the skeletal vibrations of pyridine in these solutions suggests that the electron distribution in the pyridine molecule has been altered significantly. 4. A mole ratio study Of alkali metal salt solutions in pyridine and acetone by nmr techinques should be under- taken to determine the solvation numbers of the alkali metal ions in solution. 5. Vapor pressure studies, or studies of some other colligative property, should be undertaken in pyridine and acetone to determine if there are any polymeric species in solution. 7 and Na23 salts in solution by magnetic 6. A study of Li resonance techniques might provide some evidence for the existence of polymeric species in solution. 124 571711 111411111er1 3 1293 Inunmmmummm