SOME STUDIES OF MOLECULAR COMPLEXES OF INTERHALOGEN COMPOUNDS Thai: for tho Dogm of Ph. D. MICHIGAN STATE UNIVERSITY W. Keith Meyer 1958 TH ESIS 0-169 “"3. Want... a m. dim-3‘4" 4" I. 3"" '“ IR Y {Jilimn 55¢. ;c s.) y. Umvusuy gaww- .m-ut: gm M This is to certify that the thesis entitled SOME STUDIES OF MOLECULAR COMPLEXES Date OF INTERHAIDGEN COMPOUNDS presented by w. Keith Meyer has been accepted towards fulfillment of the requirements for Ph.D. degree inChemiatry Oct . 2 , 1958 LIBRARY Michigan State University SOME STUDIES OF MOLECULAR COMPLEXES OF INTERHALOGEN COMPOUNDS By W. Keith Meyer A THESIS Submitted to the School for Advanced Graduate Studies of Michigan State University of Agriculture and.Applied Science in partial fulfillment of the requirements for the degree of DOCTOR OF PHILOSPHY Department of Chemistry 1958 C“ 1/.”1'2: V /ser~ ACKNOWLEDGMENT The author wishes to express his sincere appreciation to Professor M. T. Rogers for his guidance and assistance throughout the course of this work. He wishes to thank the Union Carbide Corpor- ation for a fellowship during the academic year 1957-58, and also the Atomic Energy Commission for a grant subsidizing a part of this research. fifitttfitfitt$tfititt ii VITA w. Keith Meyer was born December 21, 1929 at Sioux Falls, South Dakota. He attended Morningside College at Sioux City, Iowa and received the Bachelor of Science degree in 1951. He attended graduate school at the University of Kentucky for one year and then entered the military service for two years. After returning to the University of Kentucky. he received the Master of Science degree in 1955. He entered the graduate school at Michigan State University in 1955. He is a member of the American Chemical Society, Sigma Xi, Sigma Pi Sigma and Alpha Chi Sigma. iii SOME STUDIES OF MOLECULAR COMPLEXES OF INTERHALOGEN COMPOUNDS By W. Keith Meyer AN ABSTRACT Submitted to the School for Advanced Graduate Studies of Michigan State University of Agriculture and Applied Science in partial fulfillment of the requirements for the degree of DOCTOR OF PHILOSOPHY Department of Chemistry Year 1958 Approved 2 274' a; g/a ABSTRACT A series of new molecular complexes has been prepared. Complexes of iodine pentafluoride with pyridine, 2-methylpyridine, dioxane, 2-f1uoropyridine, and trifluoroacetic anhydride have been made. Complexes of IBr, 101 and ICl3 with dioxane and a variety of substituted pyridine derivatives have been made.1 Melting points, molecular formulae and qualitative observations of the colors and stabilities of most of the complexes have been tabulated. The electric moments of 3-chloropyridine, 2-f1uoropyridine, 3-f1uoropyridine and 2,6-dimethy1pyridine have been measured in benzene solution at 25° C. The electric moments of seventeen molecular complexes of organic amines and ethers with inter- halogen compounds have been measured at 25° C. in solution in a nonpolar solvent. The results were interpreted in terms a? a model in which a lone pair of electrons of the nitrogen or oxygen atom is donated to the iodine atom of the interhalogen compound. Che dative bond so formed is polar in character and the percent charge transfer was calculated for each complex from the differ- ence between the observed moment and the moment calculated for no interaction. The dative bond was assumed to be linear with the axis of the interhalogen compound and the plane of the pyridine ring except in the case of iodine pentafluoride where it was assumed to be at an angle of 30° to the axis of the tetragonal pyramidal molecule. Partial phase diagrams were completed for the systems pyridine-iodine pentafluoride and dioxane-iodine pentafluoride. There is evidence of 1:1 compound formation in both systems and, in addition, of a 1:2 compound in the dioxane-iodine penta- fluoride system. From the ultra violet absorption spectra of a series of solutions of the complexes in carbon tetrachloride the dissociation constants of five of the complexes were obtained. II. III. IV. V. VI. VII. TABLE OF CONTENTS INTRODUCTION 0 O O O O O O O O O O 0 O O 00000 C O I O O O O O O 0000000 O O O O O O O O HISTORICAL SUMMARY................... ..... .... ..... ..... THEORY O O O O O O C O . O C O O C O O O C O O C O C C O O O O O 0000000000000 O ..... . Dipole Moment Measurement ........................ cryOBGOPyOO0.0.0.0....0.0.000000000000000000000000 EXPERIMENTAL.’OOOOOOCIOOOOOO0..0.0.0....0.00.00.00.00... DiClOCtric Constants.............................. Dissociation Constants............................ 10dometric Equivalentaeeeeeeeeeeeeeeeeeeeeeeeeeeee Freezing POint Measurement........................ Preparation and Purification of Compounds......... RESULTSOOOOCOOOOO...0.00.0000000.00.0000.000000000000000 DISCUSSION...0......OCCCO00......OOOOOOOOOOOOOOOOOOOOOC. Phase Diagram StuMOBCOOOOOOOOOOIOOOOOOOO00.0.0.0. Dipole Moments.................................... SWRYCCOOOOOODOOOOI..OOOOOOOOIOOOOOIOOOO OOOOOOOOOOOOO O APPMDIXOO0.00000000000000000.0.0.0.000...OO.....O...... Appendix AOOCOCOOCO0.000......OOOOOOOOOOOOOOOOOOOC Appendix 3.00.0.0....0...000......00.00.000.000... BIBLIOGRAPHYOOOOOOOOOOOOOCOOOOOOOOO0.0000000000000000000 vii 11 ll 19 an 2:. 27 28 29 51 88 88 88 9h 96 98 99 TABLE 1. TABLE 2. TABLE 5. TABLE #. TABLE 5. TABLE 6 . TABLE 7. TABLE 8. TABLE 9. TABLE 10. TAEIE 11. TABLE 12. TABLE 13. TABLE 1“. TABLE 15. LIST OF TABLES EQUILIBRIUM CONSTANTS OF AMINE-INTERHALOGEN COMPLEXES IN CARBON TETRACHLORIDE ............... ELECTRIC MOMENTS OF SOME METAL HALIDE COMPLEXES.. HEATS OF FORMATION OF CUPRIC AND MERCURIC HALIDE COMPLEXES WITH PYRIDINEO0.00000COOOOOCOOOOOOOCCOC CALIBRATION DATA FOR THE IODINE PENTAFLUORIDE BURENEOOIOOOCOIOOOOOOOOOO ..... DOOOOOOOOOOOOIOOO. SOME PROPERTIES OF NEW MOLECULAR COMPLEXES OBTAINED IN THIS INVESTIGATION...... ............ . ANALYSES OF MOLECULAR COMPLEXES.................. DIELECTRIC CONSTANTS AND SPECIFIC VOLUMES OF THE BENZENE SOLUTIONS AT 15° c. .................... DIELECTRIC CONSTANTS ANg SPECIFIC VOLUMES OF THE BENZENE SOLUTIONS AT 25 c. eeeeeeeeoooooeeeeeee DIELECTRIC CONSTANTS AND SPECIFIC VOLUMES OF THE CARBON TETRACHLORIDE SOLUTIONS AT 25 C. ....... DIELECTRIC CONSTANTS ANB SPECIFIC VOLUMES OF THE DIOXANE SOLUTIONS AT 25 C. ................... DIELECTRIC CONSTANTS ANg SPECIFIC VOLUMES OF THE BENZENE SOLUTIONS AT 35 c. eeeeeeeeeeeeeeeeeeee DIPOLE MOMENTS, MOLAR POLARIZATIONS, MOLAR REFRACTIONS AND EMPIRICAL CONSTANTS FOR MOLECULAR COMPLEXES...................... ...... ............ CRYOSCOPIC DATA FOR PYRIDINE—IODINE PENTAFLUORIDE SOLUTIONS. ..... ........... ............. . ....... .. CRYOSCOPIC DATA FOR DIOXANE-IODINE PENTAFLUORIDE SOLUTIONSOOOOOOOOOOO0.0.0.000...00.000.000.000... ULTRAVIOLET ABSORPTION SPECTRA OF THE AMINE- HAIJOGEN COMPLHESQOOOOOOOOOO ........ 0.00.00.00... viii 10 36 49 52 57 75 78 79 82 TABLE 16. TABLE 17. TABLE 18. TABLE 19. TABLE 20. LIST OF TABLES, continued EQUILIBRIUM CONSTANTS OF AMINE-HALOGEN o COMPLEXES IN CARBON TETRACHLORIDE AT 25 C. .... OBSERVED AND CALCULATED DIPOLE MOMENTS OF MOLECULAR COMPLEXES........ ....... .. ..... . ....... OBSERVED AND CALCULATED DIPOLE MOMENTS OF SOME SUBSTITUTED PYRIDINESOCCOCOOOCOCOOOCOCOOCCOOOOCOC DIELECTRIC CONSTANTS AND SPECIFIC VogUMES OF THE CARBON TETRACHLORIDE SOLUTIONS AT 25 C. ...... FREEZING POINTS OF SOME FLUOROCARBON DERIVATIVES. 82 89 93 FIGURE 1. FIGURE 2. FIGURE 3. FIGURE #. FIGURE 5. FIGURE 6. FIGURE 7. FIGURE 8. FIGURE 9. FIGURE 10. FIGURE 11. FIGURE 12. FIGURE 13. FIGURE 1“. LIST OF FIGURES Some phase diagrams of two component systems ..... The experimental cell used for dielectric constant measurements ............................ Freezing point cell .............................. Freezing point cell and cooling bath arrangement.. Iodine pentafluoride measuring burette ........... Temperature measuring circuit .................... Dielectric constants as a function of mole fraction solute for carbon tetrachloride solutions of 2-chloropyridine.ICl, 3-bromopyridine.IBr and 2-f1u0r0pyr1dino.13r.............................. Dielectric constants as a function of mole fraction solute for carbon tetrachloride solutions of 3-chloropyridine.ICl and 3-ch10ropyridine.IBr.. Dielectric constants as a function of mole fraction solute for carbon tetrachloride solutions of 4-chloropyridine.IC1 .......................... Dielectric constants as a function of mole fraction solute for benzene solutions of 2-fluoro- pyridine.IF5, pyridineJF5 and dioxane.IF5 . . . . . . . Dielectric constants as a function of mole fraction solute for carbon tetrachloride solutions of 2-fluoropyridine.ICl and 2,6-dimsthylpyridine.. Dielectric constants as a function of mole fraction solute for dioxane solutions of dioxane. ICI and dioxaneeIBr eeeeeeeeeeeeeeeeeeeeeeeeeeeeee Dielectric constants as a function of mole fraction solute for carbon tetrachloride solutions of 3-fluoropyridine.IC1 and 3-f1uoropyridine.IBr.. Dielectric constants as a function of mole fraction solute for benzene solutions of 2-f1uoro- pyridine.IFS, dioxane.IFs and 2-methylpyrazine.IF5 21 26 33 3h 37 62 63 6h 65 66 67 68 69 FIGURE FIGURE FIGURE FIGURE FIGURE FIGURE FIGURE FIGURE FIGURE FIGURE FIGURE 15. 16. 17. 18. 19. 20. 21. 22. 25. 24. 25. LIST OF FIGURES, continued Dielectric constants as a function of mole fraction solute for benzene solutions of 2-methy1- pyrazine.IF5 and trifluoroacetic anhydride.IF5 ... Dielectric constants as a function of mole fraction solute for benzene solutions of pyridine. IF , trifluoroacetic anhydride.IF5 and 3-fluoro- py 1m. .00..............OOOOOOO0.00000IOIOOOCOOC Dielectric constants as a function of mole fraction solute for benzene solutions of 2-fluoro- pyridine and 2,6-dimethy1pyridine ................ Dielectric constants as a function of mole fraction solute for carbon tetrachloride solutions or 3-Ch10r0pyridine.IC1 eeeeeeeeeeeeeeeeeeoeeeooco Dielectric constants as a function of mole fraction solute for benzene solution of dioxane.IF5 eeeeeeeeeeeeeeeeeeeeeeeeeeeeeeeeeeeeee Phase diagram of the system pyridine-iodine Pentatluoridc ....OOOOOOOOOOOOOOOOOOO0.000.0.00... Phase diagram of the system dioxane-iodine pentafluoride ....OOOOOOOOOOOOOOOO00......0.00000. Absorption spectra of carbon tetrachloride solutions containing a constant concentration (0.00011 M) of 101 and varying amounts of 2-fluoro- pyridine.......................................... Absorption spectra of carbon tetrachloride solutions containing a constant concentration (0.00009 M) of 101 and varying amounts of Z‘Chloropyridine .00....COO-IOOOOOOOOOOOOOOOOO00.. Absorption spectra of carbon tetrachloride solutions containing a constant concentration (0.00008 M) of 1C1 and varying amounts of 5-chloropyridine ................................. Absorption spectra of carbon tetrachloride solutions containing a constant concentration (0.00005 M) of 101 and varying amounts of u-chloropyridine.................................. xi 70 71 72 73 7h 81 83 85 86 LIST OF FIGURES, continued FIGURE 26. Absorption spectra of carbon tetrachloride solutions containing a constant concentration (0.00001 M) ICl and varying amounts of Z-fluoropyridine eeeeeeeeeeeeeeeeeeeeeeeeeeeeeeeee 87 xii INTRODUCTION Molecular complexes of organic amines and ethers with the halogen and interhalogen compounds have been intensively studied in recent years. However, only a single halogen fluoride complex, dioxane-iodine pentafluoride, has been reported. Since the physical and chemical properties of halogen fluoride complexes with ethers and amines should be of interest as they might provide a new class of mild fluorinating agents, the preparation and study of a variety of such complexes was undertaken. In addition an investigation of a variety of new complexes containing 1C1, IBr and 1013 similar to the iodine pentafluoride complexes was undertaken. The stability of complexes of this type may be established by measuring dissociation constants spectrophotometrically. Infor-, mation concerning the nature of the dative bond between the two molecules of the complex may be obtained by measurement of the electric moment. A study of the electric moments of a series of molecular complexes was therefore undertaken with dissociation constants also being measured in several cases. The temperature-composition diagram for the condensed phases of a system such as ICl.pyridine shows the number of com- pounds formed, their composition, approximate stabilities and their melting points. In several cases a study of the phase diagram was undertaken to provide this information. 1 HISTORICAL SUMMARY One of the classical reviews of complexes up to the time of its publication in 1927 was Pfeiffers (38) book, Organische Molekfilverbindugen. In 195# Andrews (1‘) published a review article on aromatic molecular complexes1 limited to complexes of the "donor-acceptor" type. Work on molecular com- plexes of halogens and interhalogens with nitrogen and oxygen- containing compounds will be reviewed here. Molecular complexes are formed by the interaction of compounds in which there is a sharing of electrons between the components, resulting in a semblance of a chemical bond. One of the components donates electrons while the other acts as an electron acceptor. These interactions are weak as evidenced by their low heats of forma- tion (18, 27, 29, 52) in the order of a few kilocalories per mole as compared to normal chemical bonds with heats of formation in the order of tens of kilocalories per mole. For this reason these substances are called complexes rather than compounds. A great deal of research has been done on complexes of iodine with pyridine. Chatlet (6) reported the isolation of the 2:1 pyridine-iodine complex as yellow transparent crystals which quickly decomposed to iodine and pyridine. He also 1The terms molecular complex, addition complex, mo- lecular compound and molecular addition compound are often used with substances of this type. The author will use the first term fer description of these products. 2 reported isolation of two hydrated complexes with the composi- tional IZ.P3(H20)6 and IZ’PyA(H20)2h (7). The dipole moment of the anhydrous pyridine-iodine complex has been reported as #.5 D. in cyclohexane (30) and 4.17 D. in benzene (59). The latter value is probably in error since iodine forms a weak molecular complex with the solvent itself. Mulliken (36) suggested possible structures for the pyridine-iodine complex and Syrkin and Anisimova (59) proposed similar structures C)”; (3/: 'CK: to explain the high dipole moment of the Py.I2 complex. The work of Hassel (19-22) on complexes of oxygen and nitrogen- containing compounds with interhalogen compounds indicates that some of these proposed structures probably don't exist. He has shown that the pyridine-iodine monochloride complex is planar with the N-I-Cl group linear and that in the dioxane-iodine monochloride complex the O-I-Cl arrangement is linear. 1The symbols Py and BPy will be used for pyridine and bipyridine respectively in formulae. From thermal analysis Fialkov (IA) found two compounds in the pyridine-iodine monochloride system having mole ratios of 1:1 and 1:2 respectively. In solution in polar solvents the complexes dissociate to give Py.I+ and IClZ- ions. Popov (39) has shown that both pyridine and 2,2'-bipyridine complexes with iodine monochloride dissociate in acetonitrile to form the Ic12’ ion and Py21+ and spy:+ ions respectively. a. was unsuccessful in attempting to prepare the Py.2ICl complex. The dipole moment of the pyridine-iodine monochloride complex has been measured in benzene and found to be 8.20 D. (26). This value seems a little high in comparison to the electric moment of the pyridine-iodine complex. Williams (67) succeeded in isolating the pyridine- chlorine and the pyridine-bromine complexes, which have melting points of #70 C. and 620 C., respectively. The chlorine complex decomposes spontaneously in air. Popov and Rygg (#0) studied, by spectrophotometric methods, the molecular complexes of pyridine, 2-methylpyridine and 2,6-dimethylpyridine with iodine, iodine monochloride and iodine bromide. They calculated equilibrium constants for the dissociation of the complexes into their components in carbon tetrachloride solution at 250 C. These dissociation constants are listed in Table 1. TABLE 1 EQUILIBRIUM CONSTANTS OF AMINE-INTERHALOGEN COMPLEXES IN CARBON TETRACHLORIDE Complex Eguil. Const. Pyridine. ICl . . . . . . . . 2.07 x 10-6 -6 2-Methy1pyridine.ICl. . . . . 1.12 x 10 2,6-Dimethy1pyridine.ICl. . . 1.12 x IO"5 Pyridine.IBr. . . . . . . . . 7.73 x 10"5 2-Methylpyridine.IBr. . . . . #.25 x 10'"5 2,6-Dimethy1pyridine.IBr. . . 2.67 x IO’I Pyridine.12 . . . . . . . . . 9.88 x 10"3 2 . . . . . 6.68 x 10-3 2,6-Dimethylpyridine.12 . . . 1.97 x 10"2 2-Methylpyridine.1 Besides the complexes listed above, the only other organic complex reported with iodine bromide is the one with dioxane (#4, 5A). A slightly different approach to the problem of deter- mining whether there is interaction between halogens and the solvent is by use of infrared spectra (37). Iodine monochloride has a fundamental absorption band at 375 cm.”1 in carbon tetra- chloride. Shifts of this stretching vibration frequency to longer wavelengths are observed as the iodine monochloride complexes with the donor. This band is shifted to 270 cm."1 for the strongest complex, pyridine-iodine monochloride. The intensity of the absorption also increases as the vibration frequency shifts to smaller wave numbers. Molecular complexes of halogens and interhalogens with oxygen-containing compounds have been mainly limited to ethers and alcohols. Lilich and Presnikova (52) determined the stability constants of dioxane and methanol complexes of iodine, bromine, iodine monochloride and iodine bromide in carbon tetra- chloride solutions. Keefer and Andrews (28) also determined equilibrium constants of complexes of iodine and bromine with some aliphatic alcohols and ethers. Mulliken (35) postulated the following structures for iodine-ether complexes R‘\\“6,/CI . R‘\\\\ I O . and 0.. 3 e .1" e . I- w/ E/ where the iodine axis is perpendicular to the R-O-R plane. He proposed that the iodine-ketone complexes resonated between two structures of this type where the axis of the iodine molecule is coplanar with the ketone skeleton. Similarly, Syrkin (59) proposed structures of the type -"°..\ FM. fl)“ ..°'l-l°"0/ 2 U .",,-o/ 2 U a? for dioxane complexes with iodine, bromine and sulfuric acid. These structures were postulated to account for the dipole 7 moments, which are 0.95 D., 1.50 D., and 8.65 D., respectively, in dioxane. In the light of Hassel's work on crystal structures (19) some revision of these structures is necessary. He has shown that halogen and interhalogen complexes with ethers, thio- ethers and amines are linear with respect to the plane of the molecule and the halogen axis. Similarly in the dioxane-bromine complex, the O-Br-Br arrangement is also linear and the O-Br bond is not in the plane formed by the oxygen atom and its two ad- jacent carbon atoms. KortUm and Walz (31) reported the value 5.00 D. for the dipole moment of the dioxane-iodine complex in cyclohexane solution. Fairbrother (15a) reported the value 1.5 D. for the dipole moment of the dioxane-iodine complex in dioxane solution. Rheinboldt and Boy (AA) give the melting point of the 1:1 dioxane-iodine monochloride complex as 56-8o C., while Hassel gives 10}0 C. as the melting point of a 1:2 dioxane-iodine mono- chloride complex formed in the vapor phase. Only one iodine pentafluoride complex is reported in the literature and that is a 1:1 complex with dioxane (51). Skelly and Popov (#1) determined the electrical con- ductance of some polyhalogen complexes in acetonitrile. The polyhalogen complexes such as (CEBMNIBr2 and (CH3)4NIBrCl behaved as strong electrolytes. The electric moments of the hydrOgen halides measured in several solvents show that they form complexes with dioxane (65). Spectrophotometric studies (A) show that iodine, iodine monochloride and bromine interact to some extent with trifluoro- acetic acid. In the last several years, a considerable amount of work has been done on molecular complexes formed by metal halides with organic oxygen and nitrogen compounds. Curran and wenzke (10, 11) reported electric moments of the mercuric halides in dioxane: however, they were probably actually measuring the moment of a mercuric halide-dioxane complex. In determining the solubility of aluminum bromide in pyridine from ~10° to 70° C., Mfiller, 33.31. (54) found that several compounds were formed. I. A. Sheka, 33 9;. (53—57) and H. Ulich, 53 3;. (62, 65) have measured dipole moments of various metal halide complexes in benzene solution. These are tabulated in Table 2. The aluminum halide complexes containing two moles of dioxane (Alx3.20#B802) appear to dissociate in solution to give the monodioxane complex since the dipole moments for both complexes are identical. It has been observed that the dipyridine complex of mercuric iodide (70) dissociates in benzene solution, probably into pyridine and a monopyridine complex. The mono- pyridine complex has been isolated and the heat of formation of mercuric halide and cupric halide complexes with pyridine have been measured (2“) as listed in Table 5. ELECTRIC MOMENTS OF SOME METAL HALIDE COMPLEXES ‘1 D., Complex AlBr3.20#H802 . . “123’6'CAH802 A1013.C#H802 AlCl ’ZC4H802 . . 3 AlBr3.PhZCO e e e AlBr3.o-C7H7N02 . AlBr3.Anisole . . Amr3.Ph NH e e e 2 3 20 . . . AlBr3.HZS e e e e AlCl .Anisole . . AlBr .Ph 3 A -. . 1C130 C7H7NO2 A e "’ e 1013 p C7H7N02 AlClB.(C2H5)20 . TABLE 2 5-23 5.23 4.62 5.19 5-19 8.41 9.50 9.76 6.58 9.56 6.68 6.56 5.14 6.45 9-13 9.48 7.79 6.54 8.92 9.68 6.54 Complex A1013.C6H5Nli2 . . A1C13.C6H5N02 . A . 1C13 C6H5COC1 A1C13.(C6H5)ZCO . T1615°204H802 T1013.5C5 T1C13.20937N . . H5N BC13.(C235)20 . . BC13.CHBCN . . . BCl .C R CN . . . . . 3 2 5 BeClZ.2(C2H5)20 . BeBr2.2(CZH5)20 . InBr3.(C285)20 TiC14.CBR7CN . TiClh.C6H5CN . SnCl 1,.C 6H5CN . SnClu.ZCH COC6H 3 5 SnC14.206H5CH0 . I! D. . . 6.86 9-05 8.92 8.72 5.66 4.07 5.68 5.98 7.65 7.75 6,84 7.57 5.04 6.05 6.16 6.55 5.60 7.70 8.70 7.50 10 TABLE 5 HEATS OF FORMATION OF CUPRIC AND MERCURIC HALIDE COMPLEXES WITH PYRIDINE Complex - A H kca1./mole CuC12.2Pyridine . . . . . . . . 28.15 CuBr2.2Pyridine . . . . . . . . 24.89 CuBr2.6Pyridine . . . . . . . . 45.20 HgClZ.Pyridine . . . . . . . . 11.1u HgC12.2Pyridine . . . . . . . . 17.22 HgBr2.Pyridine . . . . . . . . 9.19 HgBr2.2Pyridine . . . . . . . . 15.u5 HgIaozpyridj-ne e e e e e e e e 15.22 The dioxane complexes of sodium, lithium and potassium iodides have been isolated (46). The alkali chlorides and bromides do not form complexes. The dioxane complexes of di- valent metal halides have been studied by several investigators (47, 68, 69). Magnesium bromide forms stable 1:2 complexes with acetone, ethyl acetate and propanal (55), and with acetonitrile a 1:1 complex is formed. Magnesium bromide does not form stable complexes with fluorinated esters, aldehydes and nitriles. Com- plexes of the tin tetrahalides have also been prepared (45). THEORY Dipole Moment Measurement Electric dipole moments are very useful in determining the structures of organic and inorganic compounds. The general theory of dipole moments and the derivation of equations per- taining to the theory are adequately described in the literature (5, 51, 58) so only a very brief outline will be presented here. When two atoms of differing electronegativities are joined together by a chemical bond, the more electronegative atom will accumulate a negative charge and the remaining atom will be more positive in character. This separation of charge by some distance constitutes an electric dipole. The dipole moment is the product of the electrical charge and the distance of separation. This electric moment (,0) is a vector quantity since it possesses both magnitude and direction. Thus in a nmlecule composed of several atoms and having several individual bond moments, the total dipole moment is the vector sum of all the bond moments. The electronic charge is 4.8 x 10"10 e.s.u., so that if a positive and negative electronic charge are sep- arated by 1 R, which is a distance of the order of a bond length, a moment of 4.8 x 10-18 e.s.u. is created. This is generally reported as 4.8 Debye units. The usual methods of obtaining the dipole moment of a substance depend upon measurement of the dielectric constant. 11 In an electric field, the electrons and nuclei are displaced from their mean positions so that both polar and non-polar molecules will become polarized. This polarization corresponds to individual moments mE and m.A induced by displacement of the electrons and nuclei by the electric field F. If an electric field of unit strength induces the moments ME and MA in a non- polar molecule, then in a field of strength F, the average moment over all molecules is: m = (ME + MA)F (1) This is often referred to as "induced" or ”distortion" polar- ization. If the molecules have a permanent dipole moment, rather than being nonpolar, they will tend to orient themselves to Oppose the field. This alignment is disturbed by thermal agita- tion of the molecules. The resulting average orientation will be at some point between the original position of the dipole and the perfectly aligned position. The result is that a slight excess of dipoles will be in Opposition to the field at any par- ticular time and this corresponds to a further moment, the "orientation" polarization 50. This orientation polarization was first evaluated by Debye (15), with the aid of certain approximations. He found' a = ”Zr (2) 5kT Where k is the Boltzmann constant and T is the absolute temper- l3 ature. The total moment (at) is then: at: (ME+MA+—L3k: )F (3) To evaluate p from experimental data, Flt/F must be replaced by another expression containing known or measurable quantities. This can be done by using the expression developed by 0. F. Mosotti (1850) and R. Clausius (1879). Their equation was obtained by a consideration of the definition of the dielectric constant and the force acting upon a polar molecule at the center of a spherical cavity in a dielectric which has been placed in an electric field. They derived the relationship: 5 t _ M e e E " 1 (a) F - Nd EW’ 6‘ + 2 where 8 is the dielectric constant, M is the molecular weight, d is the density and N is Avagodro's number. Combining Equations 5 and 4 gives: a .M 41TN(M.B+MA+// T: 3 5kT P = E. t ‘E + p..- ) (5) p as the complete expression for the total molecular polarization (Pt)' The total polarization is equal to the sum of the elec- tronic, atomic and orientation polarizations: Pt = PE + EA + PC . (6) It should be noted that from assumptions made in its derivation Equation 5 is valid only for gases at pressures where the molecules exert no mutual influence on each other. With 14 certain modifications, as will be shown later, it can be applied to dilute solutions of substances in non-polar solvents. In order to make use of Equations 5 and 6 in the cal- culation of dipole moments, PA and PE must be evaluated or eliminated. If an alternating field is used in measuring the dielectric constant then, as the frequency is increased, a point is reached where the orientation polarization starts to decrease due to the fact that the inertia of the molecules prevents them from following the reversals of the field. Thus, as the fre- quency increases, the orientation polarization is completely eliminated and at still higher frequencies, the atomic polar- ization will likewise disappear for the same reason. At frequencies of the order of the wavelength of light only the electronic polarization (PE) remains. Using Maxwell's relation- ship that for measurements carried out at the same frequency, 5': n2, and measuring the refractive index at long wavelengths, the total polarization can be rewritten as: Pt=lm=n2-1._M_ (7) n + 2 d which is identical to the molar refraction (MR) as defined by the Lorenz-Lorentz formula. For a given substance this quantity is essentially constant and independent of temperature. There- fore by determination of the dielectric constant and density of a Pure substance, Po + PE + PA can be found. From the molar refraction it is possible to estimate the electronic polarization andso obtain the orientation polarization by difference if PA is 15 known. Although MR should be determined at infinite wavelength it is customary to measure refractive indices at the wavelength of the sodium D line and use MRD in place of MRm. The sum of the atomic and electronic polarizations (PA + Pi) is often taken to be equal to MRD. Some error is introduced by including PA in this expression but PA is usually small and can therefore be disregarded, or PA may be estimated as about 10% of PE. Since the orientation polarization term is I-I'II'N/Ja , 9kT we see that 2 P - MR = 41rN/l (8) T D 9kT or solving for [J and substituting the constants 7r , N and k x1 = 0.0128 \flPT - MRD)T. (9) The preceding equation applies to gases or vapors. With some modifications of the Clausius-Mosotti relation, the equation can be applied to dilute solutions of polar solutes in nonpolar solvents. Assuming that P12 = Plf1 + szz, the molar polar- ization of the solute can be expressed as: P = 12 1 - P (10) Where P12, the molar polarization for the solution is: Pia = ‘12 ' 1 . M2‘2 * M1‘1 (11) _z______. 12 + 2 d12 Where the subscripts l, 2, and 12 refer to solvent, solute and Solution, respectively, and f is mole fraction. By measuring 16 a series of solutions containing low concentrations of solute, the corresponding P2 values are obtained. These are then plotted graphically versus f2 and the best straight line through the points extrapolated to zero concentration to give the true polar- ization (P20) of the solute in the absence of solvent effects. This value of P20 is combined with the molar refraction ‘ (P2o - MRD) and used in Equation 9 to compute the dipole moment. An error is introduced in the dipole moments measured by this method, due to assumptions in the derivation of the equations. This error is assumed usually to have a maximum value of a few tenths of a Debye unit (66). The solvent effect, which is the difference between dipole moments measured in the gas phase and in dilute solutions of nonpolar solvents, is usually not greater than 0.2 D. (60). The dipole moments in this work were calculated accord- ing to the method suggested originally by Hedestand (25) and {modified by Halverstadt and Kumler (17). These workers have shown that both dielectric constant ( 5,2) and density or specific volume (l/d12 = V12) are usually linear functions of the Inole fraction of solute for dilute solutions. The relationship of the dielectric constant and density to mole fraction is then: £12 61(1 + 3 f2) (12) (112 ‘11” +,a £2) (13) where a and fl are the slopes of the respective plots of 1? dielectric constant and density versus mole fraction of solute f2; also 61 and 6.1 are the corresponding intercepts at infinite dilution. The substitution of Equation 11 into Equation 10 results in the following expression: _P2=£l2-1._M_2_+f1M1(€12-1._1__ 51'1.__:_L_)(11+) 5_12+1 612 'r"2"' £——12+2 c112 ”31+2 d1 Substitution of Equations 12 and 15 into Equation 14, expansion of the resulting expression and simplification, gives the following equation: P2 ___ £12 - 1 . :12- + £1141 53:1 - p (:1 ;1)[51(1 em :2) + 2])(15) PF + an d12( ( £1 + 2)( £111 + a: £2) f 2')” " . . . _ (n _ _ Now in the limit as f2'-+I 0, P2 _ P2 , d12 — (11 and fl - 1. Equation 15 then becomes: p20: 3‘51’5 + "2 " PHI . £1 " 1 (16) 2 d 5' + 2 d1( £1,* 2) 1 1 In this work, 5' was calculated from plots of specific volume versus mole fraction and in this case 5' = ’P and &'= 0581 . ‘31 so that Equation 16 simplifies to: co _ 5 3' M , PO - 1 2 + (M2171 + p M1) e ( £3.+-2) .51 + 2 V1 81 " 1 (17) CPhis equation was used in all the calculations reported here. {Phe values of 5i andv1 used in the calculations were-the Eiccepted values for the pure solvent and not the extrapolated 18 values of the £i2 versus f2, and V12 versus f2, plots. The extrapolated values of' fi’and.vl, in most cases, agreed fairly well with the accepted values for pure solvent. The additivity of bond refractions was assumed in the calculations of molar refractions (MRD) and values of bond re- fractions reported by Vogel (64) were used. To determine the dielectric constants of the solutions, it was necessary to calculate a cell constant for the experi- mental cell. This was done by measuring the capacitance of the cell at the desired temperature, filled first with dry air and then pure solvent. The cell constant (k) was calculated from the following equation: k = cair ' Cs (18) Si - 1 where the values for 63., the dielectric constant of the solvent at 250 C.. were as follows: benzene, 2.2725; dioxane, 2.2080; and carbon tetrachloride 2.2280. The dielectric constant of benzene at other temperatures was given (51) by the expression: 6'. ... £25 . ( /£/,d‘t)(t - 25) (19) where (T/d‘t = -0.0019. The dielectric constant of each solu- ‘tion was calculated from the following equation using the cell «constant along with the capacitance readings for the cell filled V'ith air and then filled with a solution of a concentration (Csoln. ) ' 512 = (Cair ' Csoln.) + l (20) k 19 The density of each solution was determined at the same temperature as the dielectric constant. The slopes (1' and 5' were determined from the data by the method of least squares and used in Equation 17 to determine the molar polar- ization. The dipole moment (I!) was calculated from Equation 9. Cryoscopy The depression of the freezing point of the solvent by solute in ideal solutions is a colligative property which is dependent on the concentration of the solute in solution and not on the nature of the solute. The slight pressure dependence of freezing point is usually disregarded since most measurements are made at atmospheric pressure which varies only slightly. The depression of the freezing point ATf is related to the mole fraction f2 of the solute (if pure solvent separates as solid phase) by the equation: ATf = o 2 (21) where R is the gas constant, To is the freezing point of the pure solvent, and Lf is the molar heat of fusion of the solvent. This equation is useful in dilute solution where f2 is small and it can be shown that: ATf == Kfm (22) where m is the molal concentration of the solute and Kf the molal freezing point depression constant and is defined by the 20 equation: 2 K - RTo (23) f _ 1000 L 7E1 where M1 is the molecular weight of the solvent. For a condensed system of solid and liquid phases in equilibrium, the relationships can be expressed by a temperature- composition diagram. The curve of a temperature-composition diagram can be called either a "freezing-point depression curve" or a "solubility curve" since it illustrates both phenomena. The curve can be considered to represent the depression of the freezing point of the first substance by the second substance or the solubility curve of the second solid with the first as solute. The term "solubility" is used when the solid phase that separates is the solute, and "freezing point" when the solid is the solvent. A general discussion of the many different kinds of phase diagrams of condensed systems can be found in the liter- ature (15, 16, #3). A system of two components can show a number of different types of composition-temperature diagrams. Diagram A of Figure l is a simple system showing a single eutectic point and two freezing-point curves. Diagram B of Figure 1 represents a more complicated system in which the two components form a compound in a 1:1 ratio. This diagram also shows solid solution of the 1:1 compound in the one component. Since the variables of a condensed system are composi- tion and temperature, the methods for determination of phase diagrams can be classified as either isothermal or isoplethal. 21 u C D '- 4 C m 0. I! u ... O MOLE FRACTION l A m 1 I! i D 1 '— u < 5 m n U I 0- I I: u U .... [I / / 0 MOLE FRACTION I 8 Figure 1. Some phase diagrams of two component systems. 22 Isothermal methods involve the determination of the solubilities at known temperatures. For solid-liquid solubility curves this may be accomplished by the separation of the phases at equilibrium followed by analysis of the phases. The isolation of the liquid from the solid is often difficult, especially if one component is volatile or if the solid is very finely divided, mechanically retaining the liquid. The composition of the saturated solution may be deter- mined without separation from the solid by plotting isothermally some property of the liquid as a function of composition. Starting with an unsaturated solution the property will vary smoothly with composition until the solution is saturated at which point a break will occur. If this is performed at enough temperatures the plot of composition versus temperature can-be worked out. Isoplethal methods involve the measurement of the tem- perature of phase transition. One of the simplest of the iso- plethal methods is the visual method in which the temperature of appearance or disappearance of a particular phase is observed. This has an advantage in that it can be repeated on the same sample for checking results. For solid-liquid systems this is very readily applied to the first appearance of crystals from solution. The method of thermal analysis is applicable in those cases in which a sample is heated above the phase transition point, allowed to cool under controlled conditions, and a plot 23 made of temperature versus time. The break in the temperature- time curve corresponds to the phase transition and is due to the evolution of the heat of transition of the material. The method of dilatometry is similar to that of thermal analysis. The volume of the sample of known composition is plotted versus temperature at constant pressure. Since time is not a variable in this method, the sample may be held at a con- stant temperature to insure equilibrium in the sample. This method overcomes the question of internal equilibrium in the thermal method. EXPERIMENTAL Dielectric Constants The dielectric constants were measured by use of an apparatus employing the heterodyne beat method, in which the frequencies of two oscillating circuits are matched. One of these circuits contains a two hundred kilocycle quartz crystal which serves as the control element in the fixed frequency circuit. The other circuit contains the experimental cell, a variable precision condenser (General Radio, Type 722D), and an inductance which are connected in parallel in one of the tuned circuits of the variable frequency oscillator. The frequency )'of this second circuit is given approximately by: where L is the inductance and C the capacitance. Changes in the cell capacitance must be matched by a compensating change in the precision condenser. The oscillations from the two circuits are fed into a mixer tube (68A?) where the emerging frequency is the difference of the two. This is detected by a speaker, the pitch decreasing as the frequencies are matched. For more sensitive detection of the null point a "Magic Eye" indicator tube (635) was used. This apparatus is very similar to that described by Chien (8). 2# 25 When measuring the dielectric constant, the frequencies of the two circuits are matched with air in the cell and then a solution is placed in the cell. The change in capacitance of the cell on introducing the sample must be matched by a change in the capacitance of the precision condenser to bring the frequency back to the null condition again. From the readings of the precision condenser with air and with solution in the cell the dielectric constant of the solution may be computed with the aid of Equation 20. The experimental cell (Figure 2) consisted of three concentric nickel cylinders with the middle cylinder at high potential and shorter than the two grounded cylinders. The cylinders are separated by small Teflon spacers and the outer cylinder serves as the wall of the cell. The cell is filled with about ten milliliters of the solution to be measured and allowed to come to temperature equilibrium with the bath in which the cell is immersed. The temperature in the bath was maintained 25.00 i 0.020 C. by use of a "Fisher Electronic" relay in con- junction with a knife heater. A motor driven stirrer was used to minimize temperature gradients in the bath. In general five or six solutions of a particular solute at concentrations varying from 0.0005 to 0.02 mole fraction were prepared and measured the same day. In some cases (fluoride complexes) preparation and measurement were completed in a few hours to minimize any reaction. After the capacitance change of a solution had been measured, the solution was transferred 26 D I a ’ :::::: r" M v :‘%cE M : ///’//J% V “pr/’V‘ w’/’/ / :’%/C r : w’////‘; / lg/fll / V E v { 7 Figure 2. The experimental cell used for dielectric constant measurements: A, to circuit; B, solution flask; C, to dry air and air outlet; D, Teflon spacers. directly {48) for the buo brass i grega: 27 directly from the cell into a modified Ostwald-type pycnometer (48) for density determination. All weights were corrected for the buoyant effect occuring when weighing materials in air with brass weights. The compounds, complexes and nonpolar solvents used were prepared and purified as described in a later section. Dissociation Constants The dissociation constants were determined by a spec- trophotometric method. A Beckman DK-Z spectrophotometer was used with capped quartz cells for the measurements. The measure- ments were made at room temperature (about 25° 0.). A weighed quantity of a standard solution of halogen in carbon tetrachloride was placed in the sample cell. The spectrum was determined using carbon tetrachloride in the reference cell. Then a known quantity of amine was added to the sample cell and the spectrum redetermined. This procedure was continued until there was very little or no change in the spectrum on further addition of amine. The extinction coefficient (ac) of the complex was then calculated from the final spectrum using the equation: A = ac . b . c (25) where A was the total absorbance, b the path length in centimeters through the cell and c the concentration in moles/liter. The 28 concentration of the complex was assumed to be the same as the original concentration of halogen. From the spectra of the 1:1 complex at various concen- trations in carbon tetrachloride solution, the degree of dissociation (x) was calculated from the equation: A=axebexc+acebe0(1-X) (26) inhere ax was the extinction coefficient of the halogen at the vvave length of the complex absorption peak. The dissociation constant (k'c) was calculated from Ostwald's dilution lawl. k' = x2e (27) Ilt should be noted that k'c might not be a constant because of the neglect of activity coefficients. Iodometric Equivalent The amount of elemental halogen in the molecular complexes '88 determined experimentally in the manner described below. The calculated value was found by dividing the molecular weight 0f 'tlie complex by the number of elemental halogen atoms considered to be present in the complex. N 1S. Glasstone, Textbook 2: Physical Chemistry, D. Van °Strand Co. Inc., 19% ,' p. 955. 29 Samples of the complex weighing between two and three milliequivalents were dissolved in 10 ml. of pyridine. Six grams of C. P. potassium iodide were added and the mixture cooled in an ice bath. To the cold mixture 12 ml. of 12 N. hydrochloric acid were slowly added. Heat is evolved and pyridine-hydrochloride generally percipitates and carries with it a quantity of iodine. The mixture was titrated with 0.1 N. sodium thiosulfate. After the titration had started 10 ml. of acetone were added. The acetone would dissolve the pyridine-hydrochloride and liberate the iodine. The solution was deep red at this point and usually contained some undissolved potassium iodide. The titration was continued until a yellow color was reached, then about 20 ml. of water added. Any remaining potassium iodide wo uld dissolve and the solution would be a clear yellow color. Two ml. of starch solution were added and the titration con- tinued to the starch endpoint. A blank was determined on the in dicator and components. Freezing Point Measurements The measurement of the freezing points of pure liquids and Solutions, for determining points on a solid-liquid phase diagram, were carried out using the following apparatus. The experimental cell (Figure 3) was constructed of "Pyrex" glass having an air space between the inner and outer Wall of the cell which could be evacuated to control the rate °f Cooling of the sample. The cell was fitted with a L\ Figure 3. Wf/zl M Freezing point cell: A, stirrer; B, thermocouple leads; C, to vacuum pump. 31 fluorothene1 plug with openings for the stirrer and the thermo- couple leads. The thermocouple was constructed of Leeds and Northrup 2% gauge copper wire and constantan wire. The junction on the thermocouples were silver-soldered. The thermocouple junction in the cell was sealed through a glass tube and the emerging tip was immersed directly in the solution. By this means the chance of temperature differential existing between the sample and the measuring; device was greatly reduced. The refer- ence junction was immersed in a Dewar flask filled with crushed ice and purified water. The ice employed was frozen from purified water in a closed polyethylene container. The temperature of this cold junction was assumed to be 0.000 C. The thermocouple was cali- brated at the following check points: melting point of carbon tetrachloride, «22.900 C.; transition point of sodium sulfate heptahydrate to sodium sulfate decahydrate, 32.380 C.; and transition point of anhydrous sodium bromide to sodium bromide . o . . dlhydrate, 50.67 C. These check pOints not only serve to cali- brate the thermocouple, but the entire measuring circuit, since 1She same standard cell, galvanometer and precision potentiometer were employed throughout all the calibration and experimental 1Teflon as used throughout this thesis is the registered trade name of the E. I. DuPont de Nemours 8: Co. for the tetra- uC>roethylene polymer. Fluorothene is a trifluorochloroethylene polYmer e 2Distilled water was run through a mixed-bed ion-exchange °°lumn. A metering device on the column indicated less than one part per million of sodium chloride. 32 work. Since the measured potentials at the check points agreed to within 1 x 10-6 volt with the standard potentials for a copper-constantan thermocouple (25) no correction was applied to the measured potentials. The stirrer was constructed of nickel wire and activated by a windshield wiper motor. The cell was immersed in an unsilvered Dewar flask which contained methanol (Figure #). The bath was cooled by circu- Ilating the methanol with a centrifugal pump through a copper tube and a reservoir contained in a second Dewar flask which was cooled with Dry Ice and acetone. A stopcock in the line con- trolled the flow of the cold methanol from the reservoir and therefore the temperature in the unsilvered Dewar flask. By a combination of flow rate and pressure in the jacket of the cell, any desired cooling rate of the sample could be attained. The cooling rate of the sample was generally about one degree per minute. For handling and measuring the iodine pentafluoride with a minimum of exposure to moisture, a special burette was con- structed (Figure 5). A two milliliter Pyrex pipette tube was sCribed at 0.1 inch intervals along its entire length with a diamond point. A Teflon stopcock (Lab Crest) was sealed on the bottom of the tube. The top of the tube was fitted with a fluGrothene adapter that connected the burette to a drying tube. This tube could be replaced by an aspirator bulb for filling the burette. The calibration data of the burette with mercury 33 L U M V Figure #. Freezing point cell and cooling bath arrangement: A, circulating pump; B, valve; C, reservoir; D, freezing point cell. Figure 5. Iodine pentafluoride measuring burette: A, Teflon stopcock; B, fluorothene adapter; C, calibrated tube. 35 is given in Table h. A special tube was constructed of fluoro- thene to fit over the tip of the burette and lead into a con- tainer of iodine pentafluoride. When suction was applied at the top of the tube, iodine pentafluoride would be drawn into the burette. During this filling operation, the container of iodine pentafluoride and the bottom half of the burette were encased in a polyethylene bag. The atmosphere in the bag was flushed out with dry nitrogen to remove the water vapor in the atmosphere before the filling operation commenced. When the burette was filled, a fluorothene cap was placed over the tip and the drying tube connected on the top of the burette. Knowing the density of iodine pentafluoride at a given temperature and the volume added to the solvent, the weight of the solute can be de termined . The temperature measuring circuit (Figure 6) contained both a precision potentiometer1 and an electronic recording POtentiometer with a direct-current amplifierz. The circuit was so arranged that the thermocouple output could be measured with the precision potentiometer or the path of the cooling curve could be followed on the recorder. The potentiometer in the recorder circuit was necessary in order to supply a bucking 1A type K-Z potentiometer made by the Leeds and Northrup Company of Philadelphia, Pennsylvania was used. 2A Leeds and Northrup type 9835B direct current amplifier :as used in conjunction with a Brown electronic recorder made by he Minneapolis Honeywell Company, Minneapolis, Minnesota. 36 TABLE 4 CALIBRATION DATA FOR THE IODINE PENTAFLUORIDE BURETTE Average Weight Average weight per division: Average volume per division: Total calibrated volume: 2 .7752 milliliters O 0 1+1?“ grams Burette Reading Wt. of Mercury per Division 0-5+ 2.0759 0.4151 5+-10 2.0559 .4072 10-15 2.0929 .4186 15-20 2.0465 .4095 20’-26 2.5185 .4197 26-50 1.6459 .4110 50-55 2.0717 .4145 35-40+ 2.1230 .4246 40-45 2.0685 .4137 45-50 2.0644 .4129 50-56 2.5065 .4177 56-60 1.7228 .4507 60-65 2.0596 .4079 65-70 2.1056 .4211 70-75 2.1155 .4251 75-80 2.0949 .4189 80-85 2.1075 .4215 85-90 2.1591 .4278 0.03083 milliliters 37 Type K-2 Potentiometer I.5V Brown Recording Leeds & Northrup Potentiometer D.C. Amplifier Figure 6. Temperature-measuring circuit: T.I., thermocouple input; R.S., switch for reversing the input to the recorder and potentiometer; R , student potentiometer used to obtain the bucking voltage. 38 voltage to reduce the output of the output of the thermocouple. The output was reduced to such a magnitude that when amplified, the recorder had a sensitivity of 1.25 degrees per inch. Preparation and Purification of Compounds Benzene. This compound was purified by the following procedure. The major portion of a quantity of C. P. thiophene- free benzene was frozen and the liquor poured off. The residue was melted and the freezing operation repeated. The residue was then melted again, dried with calcium chloride, and distilled and stored over sodium. The density measured at 25° C. was 0.87336. Quinoline. This compound was purified by drying Eastman White label material over anhydrous potassium carbonate, distil- ling and collecting the middle fraction. Carbon Tetrachloride. This compound was purified by freezing a major portion of C. P. material, pouring off the liquor, melting the residue and then distilling this material from calcium hydride. The middle fraction, which distilled at 76.80 C., was collected. The density measured at 25° C. was 1.5844. 2,6-Dimethylpyridine. This compound was purified by refluxing "Eastman Yellow Label" 90% 2,6-dimethylpyridine with methyl benzene sulfonate for one hour. The mixture was cooled and the upper layer separated and distilled. The distillate was dried with calcium hydride and fractionally distilled. The 39 middle fraction distilling at 142° c. at 757 mm. was used in this work. Z-Fluorgpyridine. This compound was prepared by diazo- tization of Z-aminOpyridine in 80% fluoroboric acid at a temperature below 5° C. After standing in ice for one hour, the solution was warmed to #00 C. to complete decomposition. The solution was neutralized with sodium carbonate after cooling to 5° C. The Z-fluorOpyridine was extracted from the solution with ether. Fractional distillation of the ether layer gave a product distilling at 125° c. at 742 mm. (50). No evidence of decompo- sition was indicated after several weeks. B-Fluoropyridine. This compound was prepared from 3-aminopyridine as described for 2-fluor0pyridine. The product distilled at 106° c. Trifluoroacetic Anhydride. This compound was prepared by refluxing trifluoroacetic acid over phosphorous pentoxide and distilling from fresh phosphorous pentoxide. The material distilled at 40° 0. gyrazine. This compound was obtained from Wyandotte Chemicals Corporation, Wyandotte, Michigan, and was used without further purification. Z-Methylpyrazine. This compound was "Eastman White Label" and was used without further purification. Z-Bromopyridine. This compound was "Eastman White Label" and was used without further purification. 2-Chloropyridine. This compound was "Eastman White Label" and was used without further purification. Dioxane. This compound was purified by freezing a major portion of C. P. dioxane, decanting off the liquor, melting the residue and repeating the freezing operation. The residue was then dried with anhydrous potassium carbonate and distilled. The middle fraction distilled at 1010 C. and had a freezing point of 11.660 C. It was stored in a brown bottle over fresh sodium ribbon. Iodine Bromide. This compound was prepared by addition of bromide to iodine and warming to complete solution. The iodine bromide crystallizes on cooling. The material was purified by fractional crystallization from the melt. The melting point of the final material was 151.50 C. compared to a value of #20 C. reported in the literature (9). Iodine Trichloride. This compound was prepared by addition of iodine to condensed chlorine in a flask cooled by Dry Ice. The solid product was transferred for storage into a covered fluoro- thene beaker in a dry box (2). Iodine Monochloride. This compound was prepared by adding an equivalent amount of iodine to a flask containing liquid chlorine cooled by dry ice. The compound was fractionally crystallized from the melt. It had a melting point of 270 C. (9). Iodine Pentafluoride. This compound was obtained from a supply purified by H. Bradford Thompson(6r). Pygidine-Iodine Monochloride Complex. This complex was prepared by slow addition of a carbon tetrachloride solution of iodine monochloride to an equivalent amount of pyridine in #1 carbon tetrachloride. The complex precipitated from solution and was filtered, washed with carbon tetrachloride and air dried. The complex was pale yellow in color and had a melting point of 129-1310 C. compared to literature values of 132-1390 C. (l4,l9,40, 67 ). Iodometric equivalent for the complex was 119.9 as compared to a calculated value of 120.7. Eygidine-Iodine Bromide Complex. This complex was pre- pared in the same manner as the pyridine-iodine monochloride complex. The complex is bright yellow in color and after recrystallization from methyl alcohol had a melting point of 112-1130 C. as compared to literature value of 113-1170 C. ('4. 40, 67). Iodometric equivalent for the complex was 116.1 as compared to a calculated value of 1#2.9. Pyridine-Iodine Trichloride Complex. This complex was prepared by adding a solution of pyridine in carbon tetrachloride to the iodine trichloride solution. The complex was filtered and washed with carbon tetrachloride. The complex was bright yellow and evolved chlorine on heating. Some of the complex evidently converted to the monochloride complex on heating. It melted over a range of 182-1960 C. and was evolving a gas on melting. Iodine equivalent was 77.5 compared to a calculated value of 78.1. The complex is fairly stable at room temperature. Cuinoline-Iodine Monochloridg Complex. This complex was prepared in the same manner as the pyridine-iodine monochloride complex. It has a light cream color and has a melting point of 42 1514-1550 C. Iodometric equivalent was 195.9 as compared to a calculated value of 185.7. Quinoline-Iodine Bromide Complex. This complex was prepared in the usual manner. The complex is yellow and has a melting point of 128-1290 C. Iodometric equivalent was 168.6 as compared to a calculated value of 168.0. Quinoline-Iodine Trichloride Complex. This complex was prepared in the same manner as the pyridine-iodine trichloride complex. It has a yellow color and a melting point of about 1320 C. It evolves chlorine on heating. Iodometric equivalent was 95.1 as compared to a calculated value of 90.6. 2.6-Dimethylpyriding-Iodine Monochloride Complex. This complex was prepared in the usual manner. It has a yellow color and a melting point before and after recrystallization from methyl alcohol of 98-990 C. as compared to a literature value of 112-1150 C. 0.0). .Iodometric equivalent was 132.0 as compared to a calculated value of 158.7. 2,6-Dimethylpyridine-Iodine Bromide Complex. This complex was prepared in the usual manner. This complex was recrystallized from hot carbon tetrachloride and has an orange color. The melting point was 105-1070 C. as compared to literature value of 106-1080 C. (40). There appeared to be some decomposition at the melting point. The iodine equivalent was 156.? as compared to a calculated value of 156.9. 246-Dimethylpyridine-Iodine Trichloride Complex. This complex was prepared in the same manner as the pyridine-iodine “3 trichloride complex. It loses chlorine slowly at room temperature. It has a bright yellow color, evolves chlorine on heating and melts over the range of 90-950 C. Iodometric equivalent was 90.5 as compared to a calculated value of 85.1. Pyrazine-Iodine Monochloride Complex. This complex was prepared in the usual manner. It has a dirty yellow color and sublimes on heating. It decomposes at 1930 C. Iodine crystals appear on the sides of the container after several days indicat- ing decomposition of the complex. Pyrazine-Iodine Bromide Complex. This complex was pre- pared in the usual manner. It has an orange-brown color and sublimes on heating. It decomposes at 1560 C. 2-Methy1pyrazine-Iodine Monochloride Complex. This complex was prepared in the usual manner. It was yellow in color and decomposed to a black tar in a few days with the evolution of considerable gas. 2-Methylpyrazine-Iodine Bromide Complex. This complex was prepared in the usual manner. It has a yellow-orange color, sublimes upon heating and decomposes at 1150 C. 2-F1uoropygidine-Iodine Monochloride Complex. This complex was prepared in the usual manner. Since the complex is quite soluble in carbon tetrachloride, it was necessary to con- centrate the solution before the complex would crystallize in long yellow needles from the solution. It has a melting point of 56° C. Iodometric equivalent was 130.9 as compared to a calculated value of 129.7. 49 2-F1uor0pyridine-Iodine Bromide Complex. This complex was prepared in the usual manner. It was necessary to concentrate and cool the carbon tetrachloride solution before the complex would crystallize in orange-brown needles. It has a melting point of 48-450 C. 2-Fluorgpyridine-Iodine Trichloride Complex. This com- plex was prepared in the same manner as the pyridine-iodine trichloride complex. It precipitates readily from solution. It has a bright yellow color, evolves chlorine on heating and melts over the range of 56—650 C. Iodometric equivalent was 101.6 as compared to a calculated value of 82.5. 2-Chloropyridine-Iodine Monochloride Complex. This com- plex was prepared in the usual manner. This yellow complex was recrystallized from methanol and had a melting point of 80-820 C. Iodometric equivalent was 139.3 as compared to a calculated value of 138.0. 2-Chlorgpyridine-Iodine Bromide Complex. This complex was prepared in the usual manner. It crystallizes in orange- brown needles from methanol and has a melting point of 43-450 C. Iodometric equivalent was 163.2 as compared to a calculated value of 160.8. 3-Fluoropyridine-Iodine Monochloride Complex. This complex was prepared in the usual manner. The 3-f1uoropyridine complexes are less soluble in carbon tetrachloride than the 2-f1uoropyridine derivatives which makes their separation from solution easier. It crystallizes in yellow needles and has a V3315 nee an 4.4 A '(1 (‘1 45 melting point of 95-97° c. It tends to sublime upon heating. Iodometric equivalent was 129.7 as compared to a calculated value of 130.9. 3-Fluorqpyridine-Iodine Bromide Complex. This complex was prepared in the usual manner. It crystallizes in yellow-tan needles and has a melting point of 70-720 C. Iodometric equivalent was 15#.7 as compared to a calculated value of 152.0. 33Bromopyridine-Iodine Monochloride Complex. This com- plex was prepared in the usual manner. It has a yellow color and a melting point of 9o-92° c. 3-Brom0pyridine-Iodine Bromide Complex. This complex was prepared in the usual manner. It has a yellow-orange color and a melting point of 77-780 C. Dioxane-Iodine Monochloride Complex. This complex was prepared in the usual way. The crystals are red-yellow in color and have a melting point of 92-930 C. as compared to literature values of 56-580 C. (44). This complex is unstable and decom- poses into dark colored material after several days. 3eChloropyridine-Iodine Monochloride Complex. This complex was prepared in the usual manner. The lemon-yellow material was recrystallized from methanol and had a melting point of 56° C. 3-Chlor0pyridine-Iodine Bromide Complex. This complex was prepared in the usual manner. The yellow-orange complex was recrystallized from methanol and had a melting point of 47° C. #6 #-Chlor0pyridine-Iodine Monochloride Complex. This com- plex was prepared in the usual manner. The pale yellow complex was only moderately soluble in boiling methanol. It has a melting point of about 221t-226o C. 0n heating the complex under- goes two changes: at about 100° C., it changes to an orange- brown colored solid and at about 2050 C. it changes to a yellow colored solid.. The limited solubility and high melting point indicate a certain amount of ionic character. #-Chloropyridine-Iodine Bromide Complex. This complex was prepared in the usual manner. The complex was a yellow- brown color and was slightly soluble in methanol. It had a melting point of about 193-195° c. Dioxane-Iodine Bromide Complex. This complex was pre- pared by slow addition of the halide to pure dioxane. It was filtered and washed with a small amount of carbon tetrachloride. The complex has a dark red-brown color and has a melting point of 6#° C. compared to a literature value of 650 C. (##). It is unstable and decomposes into a dark red tar within a few days. Dioxane-Iodine Pentafluoride. This complex was prepared in a dry box by slow addition of iodine pentafluoride to pure dioxane. The complex was then washed with carbon tetrachloride and dried on a suction filter. It is white and has a melting point of ll2o C. when dropped on a hot stage. After melting it immediately decomposes giving a cloud of iodine vapor. The :nelting point compares with the value 1120 C. reported by Scott and Bunnet (51). The complex can be recrystallized from hot 47 benzene if great care is exercised to prevent water vapor from coming in contact with the solution. It can be stored without decomposition at -78° C. 93225 Iodine Pentafluoride Complexes. The following complexes were prepared by mixing equivalent amounts of the indicated second component with iodine pentafluoride: 2-methy1- pyrazine-iodine pentafluoride, trifluoroacetic anhydride-iodine pentafluoride and pyridine-iodine pentafluoride. All these complexes are liquids at room temperature. Table 5 summarizes the new molecular complexes that have been prepared and reported for the first time. Table 6 lists the carbon, hydrogen and nitrogen analyses for several of the molecular complexes as reported by the Spang Microanalytical Laboratory, Ann Arbor, Michigan. 1,2-Dichloroperfluorocyclopentane. An attempt was made to prepare l,2-dichloroperfluorocyclopentane from 1,2-dichloro- perfluorocyclopentene-l (obtained from the Booker Chemical Co.) by fluorination in the vapor phase with cobalt trifluoride. Three one-inch copper tubes were filled with cobalt trifluoride suspended on steel wool. These tubes connected in series were placed in tube furnaces, so that there was a temperature gradient between the inlet and outlet of the tube of about 150° C. The inlet temperature was about 1000 C. The material to be fluorinated was introduced through a dropping funnel into a heated flask where it vaporized and the vapor was subsequently carried by a stream of helium gas through TABLE 5 SOME PROPERTIES OF NEW MOLECULAR COMPLEXES OBTAINED IN THIS INVESTIGATION Complex Color M.p. oC. Stability at room temp. Pyridine.IC13 yellow 142.6 loses C12 Quinoline.ICl3 yellow 132 stable 2,6-Dimethy1py.ICl3 yellow 90-95 loses C12 2-Fluoropyridine.IC1 yellow 56-65 loses Cl2 Quinoline.IC1 cream 154-5 stable Pyrazine.IC1 yellow dec. 193 unstable 2-Methy1pyrazine.ICl yellow dec. 115 unstable 2-F1uor0pyridine.ICl yellow 56 stable 2-Chloropyridine.lCl yellow 80-82 stable 3-F1uoropyridine.ICl yellow 95-7 stable 3-Bromopyridine.IC1 orange 90-2 stable 3-Chloropyridine.101 yellow 56 stable 4-Chlor0pyridine.101 - yellow 22h-6 stable Quinoline.IBr yellow 128-9 stable Pyrazine.IBr orange dec. 156 stable 2-Methy1pyrazine.IBr orange dec. 115 stable 2-Fluoropyridine.IBr orange 44-5 slow decomp. 2-Chloropyridine.IBr orange 44-5 stable 3-Fluoropyridine.IBr tan 70-2 stable 3-Bromopyridine.IBr orange 77-8 stable 3-Chlor0pyridine.IBr orange 47 stable 4-Chloropyridine.IBr tan 193-5 slow decomp. 49 TABLE 6 ANALYSES OF MOLECULAR COMPLEXES Complex Calculated Reported %c %H %N %c %H mm Quinoline.IBr 32.16 2.09 4.16 32.13 2.06 4.23 Quinoline.101 37.07 2.42 4.30 36.65 2.40 4.92 2-Chloropyridine.101 21.76 1.46 5.07 21.54 1.47 5.33 2-Ch10ropyridine.IBr 18.74 1.25 4.37 18.65 1.13 4.36 Pyridine.IBr 20.99 1.76 4.89 21.23 1.84 5.05 2,6-Dimethy1py.IBr 26.77 2.88 4.46 26.64 2.87 4.44 3-F1uor0pyridine.101 23.14 1.55 5.39 23.02 1.56 5.35 2-F1uoropyridine.101 23.14 1.55 5.39 23.18 1.55 4.90 3-Ch10ropyridine.ICl 21.76 1.46 5.07 21.94 1.44 5.06 50 the fluorination tube to a cold trap where the fluorinated material was condensed. This crude material was then fraction- ated to obtain the desired product. The major portion of the material distilled at 85° c. and solidified on cooling. The melting point was 300 C. The starting material has a boiling point of 89.5-91o C. and probably is very hard to separate from the product by fractional distillation. Cooling curves of the purest material obtained by fractional crystallization and sub- Ilimation still indicated that it was not a pure compound. This xnaterial had a melting point of 34° C. The solid material can Ice sublimed into white crystals that turn into an amorphous glass on standing. RESULTS The experimentally determined dielectric constants ( (i2) and specific volumes (V12) of the solutions at the various temperatures are compiled in Tables 7 through 11. The graphical plots of dielectric constant versus mole fraction (f2) are shown in Figures 7 through 19. The slopes ¢' and 5' are listed in {Table 12 along with the molar polarization of the solute at :infinite dilution (P2a>) calculated using Equation 17. The molar refractions (MED) calculated from empirical constants (15),and 'the dipole moments (I!) obtained using Equation 9 are listed in Tab1e 12 also. The cryosc0pic data for the pyridine-iodine pentafluoride :system and the dioxane-iodine pentafluoride system are reported :in Tables 13 and 14 respectively. The graphical plots of freezing point versus mole fraction of solute are shown in Figures 20 and 21. The Spectroscopic data for the halogen-amine systems Zinvestigated are reported in Tables 15 and 16. The spectra of these systems are shown in Figures 22 through 26. 51 52 TABLE 7 DIELECTRIC CONSTANTS AND SPECIFIC ngUMES OF THE BENZENE SOLUTIONS AT 15 C. f 2 £12 v12 Dioxane-Iodine Pentafluoride Complex 0.000623 2.305 1.1309 .001861 2.331 1.1258 .003359 2 . 365 1 .1218 .003392 2.362 1.1219 .004057 2.377 1.1202 .004247 2.383 1.1198 Pyridine-Iodine Pentafluoride Complex 0.000654 2.322 1.1278 .001391 2.355 1.1239 .001775 2.363 1.1254 .003031 2.390 1.1239 .003281 2.415 1.1222 .004877 2.435 1.1204 .007088 2.471 1.1176 Trifluoroacetic Anhydride-Iodine Pentafluoride Complex 0.000372 2.294 1.1282 .000573 2.306 1.1280 .000641 2.310 1.1271 .000659 2.313 1.1266 .000992 . 2.331 1.1239 .001??? 2.347 1.1219 .003255 2.36? 1.1187 53 TABLE 7 , Continued f2 812 v12 2-Methy1pyrazine-Iodine Pentafluoride Complex 0.000781 2.328 1.1284 .002459 2.369 1.1240 .002704 2.372 1.1222 .002956 2.372 1.1228 .003603 2.400 1.1212 .006653 2.476 1.1139 .012012 2.601 1.1024 2-F1uor0pyridine-Iodine Pentafluoride Complex 0.000398 2.311 1.1293 .000773 2.331 1.1285 .000871 2.333 1.1282 .001585 2.350 1.1268 .002235 2.367 1.1256 .004394 2.463 1.1198 .012044 2.756 1.1011 54 TABLE 8 DIELECTRIC CONSTANTS AND SPECIFIC VOLUMES OF THE BENZENE SOLUTIONS AT 25° c. 1’2 £32 v12 2,6-Dimethylpyridine 0.002561 2.285 1.1442 .005421 2.299 1.1459 .008902 2.515 1.1459 .011601 2.524 1.1440 .020718 ' 2.561 1.1429 2-Fluoropyridine 0.004551 2.338 1.1424 .005176 2.344 1.1424 .008123 2.387 1.1417 .008664 2.392 1.1416 .009957 2.414 1.1411 .017703 2.525 1.1385 .027576 2.664 1.1356 3-F1uoropyridine 0.001763 2.285 1.1444 .003974 2.296 1.1437 .007126 2.314 1.1430 .010070 2.331 1.1415 .015000 2.385 1.1398 .018191 2.393 1.1383 55 TABLE 8, Continued 12 £52 v12 3-Chloropyridine 0.005280 2.303 1.1298 .011314 2.338 1.1270 .013519 2.351 1.1263 .016403 2.370 1.1248 .021110 2.396 1.1227 Pyridine-Iodine Pentafluoride Complex 0.000654 2.3040 1.1428 .001391 2.3106 1.1423 .001775 2.3421 1.1398 .003031 2.4102 1.1370 .003281 2.3914 1.1365 .004877 2.4158 1.1311 .007088 2.4988 1.1299 Z-Methylpyrazine-Iodine Pentafluoride Complex 0.0007809 .002459 .002704 .002956 -.OO56O5 .006655 .012012 2.301 2.342 2.347 2.352 2.372 2.448 2.559 1.1428 1.1382 1.1376 1.1375 1.1356 1.1282 1.1180 56 TABLE 8, Continued f2 €12 v12 Trifluoroacetic Anhydride-Iodine Pentafluoride Complex 0.001605 .002272 .002818 .003715 .004373 .004490 .005062 .005092 .009498 2.304 2.324 2.360 2.355 2.366 2.371 2.370 2.385 2.481 1.1400 1.1367 1.1339 1.1259 1.1294 1.1295 1.1273 1.1272 1.1124 Dioxane-Iodine Pentafluoride Complex 0.001295 .001782 .002372 .002426 .003452 2.299 2.309 2.320 2.322 2.343 1.1410 1.1389 1.1584 1.1385 1.1358 2-F1uoropyridine-Iodine Pentafluoride Complex 0.000398 .000773 .000871 .001585 .002235 .004394 .012044 2.310 2.330 2.333 2.349 2.367 2.462 2.756 1.1293 1.1285 1.1282 1.1268 1.1256 1.1198 1.1011 57 TABLE 9 DIELECTRIC CONSTANTS AND SPECIFIC VOLUMES OF THE CARBON TETRACHLORIDE SOLUTIONS AT 25° c. f2 £12 v12 2,6-Dimethy1pyridine 0.002478 2.239 0.6318 .003426 2.242 .6323 .004920 2.250 .6327 .009900 2.269 .6542 .010301 2.270 .6343 .019397 2.309 .6370 .022229 2.321 .6379 2-F1uoropyridine-Iodine Monochloride Complex ‘0.001557 2.268 0.6312 .003339 2.402 .6304 .004039 2.326 .6305 .006941 2.424 .6302 .008758 2.485 .6299 .012800 2.610 .6295 2-Fluoropyridine-Iodine Bromide Complex 0.000347 2.237 0.6312 .000860 2.248 .6312 .001538 2.270 .6308 .002865 2.319 .6308 .004343 2.370 .6301 58 TABLE 9, Continued I2 812 v12 2-Ch10ropyridine-Iodine Monochloride Complex 0.000251 .000534 .001606 .003157 .005334 2.237 2.256 2.315 2.401 2.517 2-Ch10ropyridine-Iodine 0.000720 .001679 .003028 .004700 .005696 .008748 2.249 2.289 2.350 2.426 2.474 2.615 0.6313 .6311 .6309 .6304 .6300 Bromide Complex 0.6311 .6307 .6304 .6296 .6294 .6280 3-F1uoropyridine-Iodine Monochloride Complex 0.000234 .000407 .000918 .001370 2.238 2.247 2.273 2.297 0.6312 .6510 ~ 6309 .6308 3-F1uoropyr1dine-Iodine Bromide Complex 0.000217 .000729 .001409 .002553 .002996 2.256 2.254 2.284 2.526 2.344 0.6312 .6510 -.6307 .6305 .6502 59 TABLE 9, Continued ‘2 512 12 3-Brom0pyridine-Iodine Bromide Complex 0.000270 .000700 .001587 .002547 ~003929 .005022 2.239 2.256 2.295 2.337 2.394 2.442 0.6314 .6311 .6305 .6500 .6298 .6286 3-Chloropyridine-Iodine Monochloride Complex 0.000443 .001153 .001456 .003480 .006070 2.249 2.289 2-303 2.403 2.515 0.6310 .6308 .6310 .6303 .6296 3-Ch10ropyridine-Iodine Bromide Complex 0.000758 .001502 .002603 .003956 2.255 2.286 2.333 2.381 0.6309 .6307 .6304 .6297 4-Ch10ropyridine-Iodine Bromide Complex 0.000174 .000410 .000492 2.234 2.242 2.247 0.6312 .6312 .6310 60 TABLE 10 DIELECTRIC CONSTANTS AND SPECIFIC VOLUMES OF THE DIOXANE SOLUTIONS AT 25° C. f2 512 v12 Dioxane-Iodine Monochloride Complex 0.000523 2.216 0.9727 .000848 2.226 .9721 .001874 2.250 .9706 .003639 2.288 .9680 .006190 2.348 .9642 .007866 2.448 .9622 .011257 2.463 .9575 Dioxane-Iodine Bromide Complex 0.001624 2.222 0.9706 .002333 2.236 .9692 .003434 2.244 .9676 .004958 2.261 .9653 .006899 2.285 .9624 .010582 2.321 .9565 .015201 2.374 .9492 61 TABLE 11 DIELECTRIC CONSTANTS AND SPECIFIC VOLUMES OF THE BENZENE SOLUTIONS AT 55° C. f2 ‘52 v12 Dioxane-Iodine Pentaflnoride Complex 0.000104 2.271 1.1556 .002397 2.299 1.1526 .004619 2.340 1.1471 .005080 2.347 1.1492 .010310 2.444 1.1380 Exefigssofioflsd e one ..mH. 2.32.82.on c 43633.08 10.3301“ 0 no engages 03.330953 39.30 sou 330» 8306.5 Odom 00 83056 e no 38300 35833 .... 053E w 62 inrkwoo ommo o mmoo o. . . \.. 1 00nd . 1 00¢.N w 4 Good 1 oomd 63 ..BH. 05303090308 0 and 3H.0§Ehnon0300n o No 30330» 0020.33.33 0350 you 3300 8.3020 0.3: uo 003008 e an 303800 3.30363 .m 0.3%: w mh00.0 — Omnwbd mwmod looms Booed w .83 I 000d ..8H553hhn 5.339% 0 H0 3035.00 93.330053 500.30 you 35.3. 00300.3 50- .«0 003.003 0 00 $552.00 3.30303 .0 0.3.8.: . .0 000.0 0 02.000 Ebb EDP 1 08.0 0 1 00nd 65 .mhuicdnoqv o «.3 £70523 0 .mhn‘équhnaoaaom 0 no 32333 0333 ..ou 35.no- aofiuauu odo- mo .8335 u no .3338 3.30303 .3 0.33» w Mwod E 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(I p P‘2 MRD p 2,6-Dimethy1pyridine col1+ 25° 9.16 0.329 101.87 33.31 1.83 2,6-Dimethylpyridine c636 25° 4.69 -o.095 100.28 33.31 1.81 Z-Fluoropyridine C636 25° 15.61 -0.194 259.24 23.85 3.39 B-Fluoropyridine C636 25° 5.67 -o.315 109.11 23.85 2.09 B-Chloropyridine C6H6 25° 5.86 -o.9u8 102.81 28.88 1.90 Pyridine-Iodine Pentafluoride c636 15° 29.55' -1.713 915.05 93.25 9.19 25° 30.99 -2.2u6 505.71 93.29 h.56 2-Flu0ropyridine-I0dine Monochloride O CCI“ 25 30.35 -O.127 555.87 #3.65 #.90 Z-Fluoropyridine-Iodine Bromide O CCl# 25 33.26 -0.278 585.07 #6.55 5.13 76 TABLE 1 2, continued Solvent '1' .OC . ‘3' ,8' P2 MRD p 2-Fluoropyridine-Iodine Pentafluoride C6H6 15° 38.29 -2.932 601.09 92.95 5.13 O 25 35.93 -2.486 578.84 92.95 5.12 2-Ch10r0pyridine-Iodine Monochloride O CClh 25 5#.36 -0.236 926.04 48.68 6.55 2-Chloropyridine-Iodine Bromide C01“ 25° 99.85 -o.379 773.03 51.58 5.99 B-Fluoropyridine-Iodine Monochloride cc1# 25° 51.19 -0.258 869.56 93.65 6.35 B-Fluoropyridine-Iodine Bromide O CClh 25 58.79 0.335 672.98 96.55 5.55 3—Chloropyridine-Iodine Monochloride O C01,+ 25 50.90 -0.194 863.38 98.68 6.31 3-Chloropyridine—Iodine Bromide ..O 0C1“ 25 40.39 -O.550 702.29 51.58 5.69 B-Bromopyridine—Iodine Bromide col“ 25° 92.66 -o.5oo 739.85 59.98 5.79 77 TABLE 12,continued Solvent '1‘ . °c . c. p’ 192CID MRD ll Dioxane-Iodine Monochloride c.31802 25° 22.83 -1.920 365.87 91.92 3.98 Dioxane-Iodine Bromide c43802 25° 10.95 -1.567 201.89 99.82 2.77 Dioxane-Iodine Pentafluoride 06H6 15° 20.75 .2.870 336.66 90.85 3.73 25° 20.06 -2.650 338.59 90.85 3.81 35° 18.51 -1.980 338.08 90.85 3.87 Trifluoroacetic Anhydride-Iodine Pentafluoride 0636 15° 22.59 -3.310 393.18 90.89 9.08 25° 21.15 -3.320 379.99 90.89 9.07 2-Methylpyrazine-Iodine Pentafluoride C6H6 15° 26.81 -2.560 933.15 99.16 9.26 25° 25.95 -2.530 933.79 99.16 9.33 78 TABLE 13 CRYOSCOPIC DATA FOR PYRIDINE-IODINE PENTAFLUORIDE SOLUTIONS O weight Weight Mole Percent F.P. c. CsflsN’(gm.) IF5 (gm.) IF5 2.1310 0.000 00.00 -91.39 2.1310 .315 5.00 -96.67 1.5702 .385 8.91 -97.89 2.1310 .622 9.91 -5o.12 1.5702 1.055 19.31 -39.69 1.5702 1.755 28.97 -15.38 1.5702 3.108 91.36 2.13 1.5702 3.778 96.33 6.83 100.00 9.60 79 TABLE 14 CRYOSCIPIC DATA FOR DIOXANE-IODINE PENTAFLUORIDE SOLUTIONS Weight Weight Mole Percent F.P. ChH802(gm.) IF5(gm.) IF5 4.8991 0.000 0.00 11.66 4.8991 1.125 8.54 6.57 2.7009 .652 8.48 6.97 2.7009 1.410 17.18 5.41 4.8991 2.708 18.04 5.12 2.7009 2.518 25.51 2.86 1-1 Ratio Compound 50.00 112.00 1-2 Ratio Compound 66.67 90.00 100.00 9.60 IO olO TC. -20 -40 -5O 80 l l 1 l O .2 .4 .6 .8 U0 MOLE FRACTION 0F IODINE PENTAFLUORIDE Figure 20. Phase diagram of the system pyridine- iodine pentafluoride: Author's conception of the remaining portion of the diagram shown by dotted lines. T ’C. 81 I20 _ IOO¢ 80 6!) 40 201 l l p l O .2 .4 6 .8 L0 MOLE FRACTION 0F IODINE PENTAFLUORIDE Figure 21. Phase diagram of the system dioxane- iodine pentafluoride: Author's conception of the remaining portion of the diagram shown by dotted lines. 82 TABLE 15 ULTRAVIOLET ABSORPTION SPECTRA OF THE AMINE-HALOGEN COMPLEXES 1 Isosbestic Complex Wave Length of ac Point Abs. Max. my my 2-Chloropyridine.IC1 556 217 596 5-Chloropyridine.IC1 502 727 592 4-Ch10ropyridine.ICl 500 604 589 2-Fluor0pyridine.IC1 555 246 597 2-F1uoropyridine.I2 440 700 484 TABLE 16 EQUILIBRIUM CONSTANTS OF AMINE-HALOGEg COMPLEXES IN CARBON TETRACHLORIDE AT 25 C. Complex Equil. 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DISCUSSION Phase Diagram Studies The phase diagram of the pyridine-iodine pentafluoride system (Figure 20) is that of a two-component system with com- pound formation at 0.5 mole fraction iodine pentafluoride. The first eutectic point is at 0.115 mole fraction iodine penta- fluoride. The l:l compound has a melting point of about 90 C. It is not a very stable compound as indicated by the change of slope of the freezing point curve at the maximum. The phase diagram of the dioxane-iodine pentafluoride system (Figure 21) appears to be that of a two-component system with two compounds formed at 0.5 and 0.66 mole fraction iodine pentafluoride. The first eutectic point is at 0.245 mole fraction iodine pentafluoride. The 1:2 compound may exhibit a true maxi- mum or it may show only a peritectic point. It exhibited its instability by dissociating into the 1:1 compound and iodine pentafluoride when dried at reduced pressure. Dipole Moments The observed and calculated dipole moments of the molecular complexes are listed in Table 17. The calculation of the dipole moment of a complex containing two group moments m1 and ma, in fixed positions was made by use of the following 88 89 TABLE 17 OBSERVED AND CALCULATED DIPOLE MOMEITS OF MOLECULAR COMPLEXES Complex Observed Calculated % Charge Shiftl 2-Fluoropyridine.IC1 4.90 n.64 2 5-Fluoropyridine.IC1 6.35 3.42 27 2-Chloropyridine.ICl 6.55 4.68 17 3-Ch10ropyridine.ICl 6.51 5.#5 26 Dioxane.IC1 3.98 1.49 20 2-Fluor0pyridine.IBr 5.15 h.36 7 5—Fluor0pyridine.IBr 5.55 5.14 22 2-Chloropyridine.IBr 5.9# #.#0 14 5-Chloropyridine.IBr 5.6% 3.15 23 k-Chloropyridine.IBr 5.79 5.15 24 Dioxane.IBr 1.65 1.21 h Pyridine.IF5 4.56 #.25 5 2—Fluoropyridine.IF5 5.12 5.15 0 2-Methylpyrazine.IF5 #.55 2.11 20 Dioxane.IF5 5.81 2.18 15 Trifluoroacetic anhydride.IF 4.07 5 1See text for explanation. equation: ll: (mla + mag + 2m1m2 cos O)h (28) where 0 is the angle between the fixed moments. The iodine mono- chloride and iodine bromide complexes were assumed to be linear. The iodine pentafluoride complexes were assumed to be at an angle of 300 with respect to the other moment. F' 0 51%” The increases in the observed dipole moments of the com- plexes over the calculated moments is due to shift of electrons between either the nitrogen or oxygen atom and the iodine atom. Possible structures of 3-f1uoropyridine.ICl that could contribute to this increase in dipole moment are: 4km -16I ’4—6: / F (I) (II) F (III) The "inner" complex of Mulliken (35) could also contribute to the dipole moment. Of these various structures, the inner complex can be eliminated. Mulliken and Reed (36) state that in the polar sol- vent pyridine, the pyridine-iodine complex has little tendency to form an inner complex. In a nonpolar solvent the tendency should be even smaller. Structure II can also be eliminated as the 91 principal structure contributing to the dipole moment. From spectrophotometric studies there is only a single absorption band in the ultraviolet region above 290 m . This hand must correspond to the complex in either structure I or III since it corresponds to the I-Cl absorption band. In structure III, the bond between the iodine and chlorine atoms would be lengthened from its normal covalent distance of 2.32 R. and the bond between the nitrogen and iodine atoms would be shorter than the normal covalent distance of 2.07 R. Hassel (19, 20, 21) has measured bond distances in the following complexes in the solid state: pyridine.ICl, dioxane.2101, dioxane.Br2, trimethylamine.ICl and trimethylamine.Iaf He has found that the halogen-halogen bond distance is the same or only slightly longer than the normal covalent bond distance, while the oxygen-halogen or nitrogen-halogen bond distance is considerably longer than the covalent bond distance. The preceding evidence strongly indicates that structure I is the predominant structure for the complex. On this basis, then,it is possible to calculate the amount of electronic charge shift from the nitrogen atom to the iodine atom that would account for the increase in the dipole moment. The difference between the observed and calculated dipole moments for 3-fluoropyridine.ICl is 2.93 D. Assuming that the N-I distance in 3-fluor0pyridine.ICl is 2.30 R. as measured by Hassel (19), this increased moment corresponds to an electronic charge shift of 22%. 92 The electronic charge shift for the remaining complexes have been calculated assuming N—I and O-I bond distances of 2.30 X. and 2.60 3., respectively, and are listed in the right hand column of Table 17. The smaller extent of charge transfer for the pyridines substituted in the two position may be due to two factors. The substituent in the two position may sterically hinder the group on the nitrogen. This could cause the two dipoles of the complex to be at an angle rather than linear as assumed in the calculation. The second factor could be that the substituent groups are electron withdrawing. They could reduce the availability of the unshared pair of electrons on the nitrogen for complex formation. The decrease in the percent of charge shift in going from the chloro to the fluoro substituted complexes can be accounted for by the increase in electronegativity of the halogen. This same trend is noticed in the iodine bromide complexes with pyridine substituted in the three position where the percent charge shift increases from 22 to 23 to 24 in going from fluorine to chlorine to bromine. A11 dipole moments of complexes were computed assuming no dissociation of the complex. The resulting errors in the moments of the complexes were estimated where dissociation constants were available. In all cases studied here k'c,where known9was 2 x 10-# or less and the resulting error in the moment of the complex was too small to be significant. The observed and calculated dipole moments of some sub- stituted pyridine compounds are listed in Table 18. 93 TABLE 18 OBSERVED AND CALCULATED DIPOLE MOMENTS OF SOME SUBSTITUTED PYRIDINES Compound Observed Calculatedl 2-F1uoropyridine 3.39 3.15 3-Fluoropyridine 2.04 1.93 3-Chlor0pyridine 1.90 1.94 2,6-Dimethylpyridine 1.81 1.80 The agreement between observed and calculated dipole moments is fairly good. 1Calculated from bond moments listed in Smyth (58). SUMMARY A series of new molecular complexes has been prepared. Complexes of iodine pentafluoride with pyridine, Z-methylpyridine, dioxane, Z-tluoropyridine. and trifluoroacetic anhydride have been made. Complexes of IBr, 161 and ICl3 with dioxane and a variety of substituted pyridine derivatives have been made. Melting points, molecular formulae and qualitative observations of the colors and stabilities of most of the complexes have been tabulated. The electric moments of 3-chloropyridine, 2-f1uoropyridine, 3-fluoropyridine and 2,6-dimethylpyridine have been measured in benzene solution at 25° C. The electric moments of seventeen molecular complexes of organic amines and ethers with inter- halogen compounds have been measured at 25° C. in solution in a nonpolar solvent. The results were interpreted in terms of a model in which a lone pair of electrons of the nitrogen or oxygen atom is donated to the iodine atom of the interhalogen compound. The dative bond so formed is polar in character and the percent charge transfer was calculated for each complex from the differ- ence between the observed moment and the moment calculated for no interaction. The dative bond was assumed to be linear with the axis of the interhalogen compound and the plane of the pyridine ring except in the case of iodine pentafluoride where 9“ 95 it was assumed to be at an angle of 30° to the axis of the tetragonal pyramidal molecule. Partial phase diagrams were completed for the systems pyridine-iodine pentafluoride and dioxane-iodine pentafluoride. There is evidence of 1:1 compound formation in both systems and, in addition, of a 1:2 compound in the dioxane-iodine penta- fluoride system. From the ultra violet absorption spectra of a series of solutions of the complexes in carbon tetrachloride the dissociation constants of five of the complexes were obtained. APPENDIX A A molecular complex of mercuric iodide with pyridine (HgIz.chhsN) was prepared. This white crystalline complex had a melting point of 100-1010 C. The dipole moment was measured in benzene solution. Zapolskii (70) found,from freezing point data in benzene, that the molecular weight was half of the expected value. The complex probably dissociates into pyridine and a monopyridine- mercuric iodide complex. On this assumption the total polar- ization of the monopyridine complex (P3) was calculated from the equation PT = Plfl + PZfZ + P3f3 (28) which is an extension of Equation 10 for a three component system. The total polarization Pk was calculated from p1,: 5:2'1 . "1‘1”‘2‘2Hi; (29) 512*2 ‘112 which is an extension of Equation 11. The dipole moment found by use of these equations was 5.88 D. The value of the molar refraction of the monopyridine complex employed (66.2 cc.) was obtained by adding the molar re- fraction of pyridine and mercuric iodide. The experimentally determined dielectric constants and specific volumes of the solutions are listed in Table 19. 96 97 TABLE 19 DIELECTRIC CONSTANTS AND SPECIFIC VOLUMESOOF THE CARBON TETRACHLORIDE SOLUTIONS AT 25 C. 2 12 12 Mercuric-Iodide Pyridine Complex 0.00053 2.2530 0.6309 .00111 2.2858 .6307 .00199 2.3311 .6301 .00390 2.k4#7 .6279 APPENDIX B The freezing points and cooling curves of some fluoro- carbon derivatives were determined in connection with the cry- oscopic work. The compounds were obtained from R. D. Dresdner of the University of Florida. The compounds and observed freez- ing points are listed in Table 20. TABLE 20 FREEZING POINTS OF SOME FLUOROCARBON DERIVATIVES Compound Freezing Point (cr3)zncoocn3 -6h.2° c. (ca3)2noccr3 -62.3° c. (c2r5)2334 -87.0° C. chrgsr6 -119.o° c. (02F5)3N Thickens to a glass -135 to -150 C. BIBLIOGRAPHY l. L. J. Andrews, Chem. Reviews, 22, 713 (l95h). 2. H. S. Booth and W. C. 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