PEWAMETHYLENWQLE COMPLEXES Win-=2 TRANSETIQN METAL HALEEIE3 meets for filo Deg?“ cf M. 5.. MICHLGAN STRTE UNIVERSE” Deiores Maureen Bowers 1968 1-H E61: 1;", &~ -. g“ LIBRARY if, Umversxty —_—— *3 ' .fl' ‘1 - ‘\ ‘ Ammgan 3:3 1:. L ‘0 i 295 I‘ M draw"! ’I i ll ABSTRACT PENTAMETHYLENETETRAZOLE COMPLEXES WITH TRANSITION METAL HALIDES by Delores Maureen Bowers This work is a part of an investigation of the donor properties of 1,5-substituted tetrazoles, in particular, of pentamethylenetetrazole (C6H10N4). The structure and the pro- perties of metal complexes of PMT are heavily dependent on the nature of the anions of the metal salts. With first row transi- tion metal perchlorates, complexes of the type MII(PMT)6(C104)2 were obtained by D'Itri (1), where PMT acts as a unidentate ligand. SUbstitution of the perchlorate anion by the halides, chloride and bromide, yields two types of complexes having the general formulas, MII(PMT)1X2 and MII(PMW)2X2. In general the solid compounds were prepared by reacting the ligand and the corresponding transition metal halide in methanol -2, 2'-dimethoxypropane mixtures. The compounds were obtained as microcrystalline powders quite stable at room temperature and insoluble in water and in most organic solvents. Magnetic suscepti- bility measurements were obtained and indicate that in all cases the complexes are spin-free. These data, in combination with reflectance spectral measurements as well as infrared spectra in Delores Maureen Bowers the 3000-100 cm-l spectral regions indicate that the complexes have an octahedral or a tetrahedral configuration. The halide complexes seem to be polymeric and contain halogen bridges. In contrast to the perchlorates, halide complexes also seem to have bridging tetrazole rings. (1) D'Itri, F. M., Masters Thesis, Mich. State Univ., East Lansing, Michigan, 1966. PENTAMETHYLENETETRAZOLE COMPLEXES WITH TRANSITION METAL HALIDES By Delores Maureen Bowers A THESIS Submitted to Michigan State University in partial fulfillment of the requirements for the degree of MASTERS OF SCIENCE Department of Chemistry 1968 @LT/#JO ACKNOWLEDGEMENTS The author wishes to express her sincere appreciation to Professor Alexander I. POpOV‘fOT his guidance, assistance, and encouragement during the course'of this investigation. A note of thanks is also extended to Dr. Norman Skelley .of The Dow Chemical Company, Midland, Michigan, for the use of the Beckman DK-Z spectrometer with reflectance attachment. Special thanks goes to my husband, Jim, and my many friends for their encouragement and helpful discussions. Finally, the author wishes to thank the National Institutes of Health of the Department of Health, Education, and Welfare, for financial aid. ii I. II. III. References . Appendix I . . . . . . . . . . . . . . . . ... . TABLE OF CONTENTS HISTORICAL . . . . . . . . . . . . . . . . . . . . . EXPERIMENTAL PART . . . . . . . . . . . . . . Reagents . . . . . . . . . . . . . . . . . . . Analytical Methods . . . . . . . . . . . . . . . . Instrumentation . . . . . . . Preparation of Coordination Compounds of. PMT . . . Dichlorobis(pentamethylenetetrazo1e)chromium(II). Dichloro(pentamethylenetetrazo1e)manganese(II). Dichloro(pentamethylenetetrazo1e)iron(II), -coba1t(II), -nicke1(II), and -copper(II). . Dichlorobis(pentamethylenetetrazo1e)zincCII). . Dibromo(pentamethylenetetrazole)manganese(II) Dibromobis(pentamethylenetetrazole)coba1t(II) and -zinc(II). . . . . . . . D1bromo(pentamethylenetetrazo1e)nickel(II) Dibromo(pentamethylenetetrazoleacopperCII). Pentamethylenetetrazole Complexes witn Chromium(II) Chloride, IronCIII) Chloride, and IronCII) Bromide. . . . . . . . . . . . . Experimental Studies. . . Solubility and ThermostabiIity. . . . . . . . . . Mole-Ratio Studies and Formation Constant Determination. . . . Composition Study in Niekel(II) Chloride Systems. Reflectance Spectra . .,. . .x. . . . . . . . . . Magnetic Susceptibility . . . . . . . Infrared Spectra of Pentamethylenetetrazole and Its Complexes with Transition Metal Halides . DISCUSSION AND RESULTS . . . . . . . . . . . . . . . . iii Page 10 11 12 14 15 15 16 16 16 16 16 17 17 19 20 20 22 28 37 . 47-50 . 51-56 TABLE II. III. IV. VI. LIST OF TABLES Page Analyses and Some Properties of Transition Metal Halide Complexes with Pentamethylenetetrazole. . . . . . 18 Composition Study of PMT-Nickel(II) Chloride Systems . . 21 Reflectance Spectra. . . . . . . . . . . . . . . . . . . 23-24 Magnetic Susceptibility Data . . . . . . . . . . . . . . 29 Molar Susceptibility Corrections for Diamagnetic Behavior . . . . . . . . . . . . . . . . . . . . .-. . . 31 Metal-Ligand and Metal-Halide Vibrational Frequencies. . 44 iv FIGURE LIST OF FIGURES Page Reaction Vessel for Chromium(II) Acetate . . . . . . . 13 Far-Infrared Absorption Spectra of Chloride conlplexeSO o o o o o o o o o o o o o o o o o o o o o 0 33-34 Far-Infrared Absorption Spectra of Bromide Complexes. . . . . . . . . . . . . . . . . . . . . . . 35-36 Dichloro(1,2,4-triazole)copper(II) Complex Structural Unit . . . . . . . . . . . . . . . . . . . 47 Figure LIST Infrared Absorption Infrared Absorption Infrared Absorption Infrared Absorption Infrared Absorption OF APPENDIX I Spectrum of PMT . . Spectrum of Mn(PMT)1C1 Spectrum of Mn(PMT)lBr Spectrum of Zn(PMT)2Cl vi 2 . 2 . 2 . Spectrum of Zn(PMT)zBr2 . Page 52 53 54 SS 56 I. HISTORICAL Tetrazole, a five membered heterocyclic ring compound containing one carbon, four nitrogens, and two double bonds, exists in two tautomeric forms; I and II (1,2) with 97% of the equilibrium mixture in form I (3). H\\\ ///H PL\\ 5 C.———-N 1 5 C ==== N 1 // \ / \ N N 2 4 N 2 \N/ \N/ \H 3 3 I II The two hydrogen atoms can be readily replaced by aliphatic or aromatic groups giving a series of 1-substituted, 5-substituted, or 1,5-substituted tetrazoles. An example of the latter type of compound is pentamethylene- tetrazole (hereafter referred to as PMT), (structure III), where the 8 ‘//’CH2 7 (I3112 \CIJHZ 9 6 ‘THZ THZ 10 5 C N 1 u l 4 N N 2 \ N/ 3 III 2 polymethylene chain is fused to the tetrazole ring forming a bicyclic compound. The first report on the synthesis and some of the pr0perties of PMT was established by Schmidt in 1925 (4). In 1928, Knoll improved upon Schmidt's method and patented a procedure (5) which is similar to that presently used by Knoll Pharmaceutical Company (6). Knoll Ltd., in England (7), has patented a process for the manufacture of PMT which is used most often today. This synthesis consists of treating cyclo- hexanone with hydrazoic acid causing ring expansion and formation of the bicyclic compound, PMT. Since PMT has been the subject of investigation in our laboratories for several years, a detailed discussion of its pr0perties is given in previous theses (8, 9, 10, 11). This work is concerned with the complexing ability of PMT, therefore a brief summary of previous studies of this prOperty is given below. The ability of PMT to act as a ligand toward transition metal ions was first reported by Schmidt (4), who obtained a PMT-mercury(II) chloride complex by mixing aqueous solutions of the two components. The composition of this metal complex was given as HgClz-Z PMT. Zwikker in 1934 (12) 'reported addition compounds of PMT formed with tetrahydrogenhexacyano- ferrate(II), trihydrogenhexacyanoferrate(III), and some salts of cadmium(II), mercury(II), copper(I), and zinc(II). These complex forming reactions were used as spot tests for PMT and the most sensitive of these Spot tests was the addition compound formed with the copper(I) chloride. The compound precipitated, PMT°2 CuCl, was very insoluble and this reaction was used for years as a quantitative means of determining PMT. Several other insoluble‘metal—PMT complexes with metal halide, perchlorate, sulfate, nitrate, and thiocyanate salts were reported by Dister (13) and by Rheinboldt and Stettiner (14). In a more detailed study of the electron donor prOperties of PMT, Popov, Bisi, and Craft determined spectrophotometrically the forma- tion constants of 1:1 PMT complexes with iodine monochloride, iodine monobromide, and iodine in carbon tetrachloride solution (15). Vaughn, g£_21,, (16) extend this study by determining the formation constants of the complexes of 7-methyl, 8-gg_o-buty1, and 8.-_t_-butyl PMT with iodine monochloride. The formation constants of the three latter complexes were feund to be slightly stronger than the corresponding complex of the unsubstituted PMT. The donor properties of PMT were determined to be moderately strong by studying the vibrational spectrum of iodine monochloride, since Person, 33.5l3, (17) have shown that the I-Cl fundamental stretching frequency was very sensitive to the strength of the interaction between the interhalogen and the donor molecule with which it was complexed. However, a solid crystalline complex could only be obtained in the PMT-1C1 system. This complex was isolated, recrystallized from chloroform and studied crystallographically by Baenziger, 33 31,, (18). It was found that PMT acts as a monodentate ligand and complexes through the nitrogen atom in the 4-position. Since silver(l) complexes with PMT had been reported by Dister (l3), Rheinboldt and Stettiner (14), and Zwikker (12), a potentiometric study of the stability constants of silver(1) complexes was conducted by Papov and Holm.(19). It was found that the formation constants for these complexes were of the order of 102. (Again only one solid complex with PMT, bis(PMT)si1ver(I) nitrate, was isolated in a crystalline form. 4 A study of divalent transition metal ion complexes with PMT was begun by POpov and D'Itri using the perchlorate salts (8, 20, 21). The investigation led to the preparation and characterization of a series of PMT complexes having the general formula MII(PMT)6(C10 where the 4),, metal ion, MII, was manganese(II), iron(II), coba1t(II), nickel(II), copper(II), and zinc(II). It was shown that the complexes had octahedral or distorted octahedral configurations and that they were soluble in water and in a number of polar nonaqueous solvents. Magnetic suscepti- bility measurements indicated that these complexes were of the high spin type. In addition to the six coordinated PMT complexes,tetrakis(PMT)- c0pper(II) and bis(PMT)copper(I) perchlorate were prepared and character- ized. Only the tetrakis(PMT)c0pper(II) and the hexakis(PMT)nickel(II) perchlorate complexes could be prepared as relatively large crystals; all the others were obtained only as microcrystalline powders. These initial studies were carried out with the perchlorates since it was desirable to minimize the possibility of interference from the anion and the formation of mixed complexes. Pentamethylenetetrazole does not seem to have very strong donor properties and the more polar- izable anions, such as halides, may well compete successfu11y with PMT for the coordination positions around the metal ion. Several transition metal halide complexes were reported by Rheinboldt and Stettiner (14) to have the following compositions: ch12-2 PMT ZnBrZ-Z PMT Zn12-2 PMT CdC12'2 PMT CdBrz'PMT CdIZ°PMT HgClz-PMT HgBrz'PMT Hg(CN)2-PMT Hg(N03)2°2 PMT CuC12-PMT AgNos-PMT S It should be noted that in the case of silver(I) nitrate P0pov and Holm (19), using the same experimental procedure, reported the formation of bis(PMT)silver(I) nitrate rather than the 1:1 complex reported above. Preparations of the PMT complexes, listed above, with zinc(II) chloride, zinc(II) bromide and c0pper(II) chloride were very similar. In each case about equal amounts of halide salt and PMT were mixed in an aqueous solution and acidified with a drOp of 2 5 solution of the respective halogen acid. The solutions were allowed to evaporate at room temperature or were concentrated under vacuum until crystalline deposits were formed. Zinc(II) complexes were first obtained in the form of oils which had to be chilled in an ice—water bath before crystals were formed. The crystals, in the form of colorless prisms for the zinc complexes and fine green needles for the c0pper complex, were filtered, washed with a few drops of the dilute halogen acid, and dried in a dessicator. The products were then recrystallized in 2 ml of water containing one drOp of the respective dilute halogen acid. Analysis for nitrogen and the halogen in the two zinc complexes corresponded within experimental error to the compositions ZnClz-Z PMT and ZnBrzoZ PMT with the respective melting points of 112° and 154-1S6°. Analysis for nitrogen, halogen, and copper indicated a composition of CuClZ-PMT with a decomposition point of about 260°. Transition metal halides are known to form mixed complexes with a variety of other heterocyclic ligands such as pyridine, triazoles, and other heterocyclic amines (23-45). These complexes take a variety of different configurations which shall be discussed below: 1. Complexes of the type L4MX2 have been reported by several investigators (22-31). Some examples are the pyridine complexes with coba1t(II) chloride and bromide and nickel(II) chloride, bromide, and iodide (27). These complexes have an octahedral configuration. 2. The complex ions of the type LMXS- have been reported by Bradbury, 3311;, (32) as atetraethylammonium salt. The organic ligand is pyridine while M is coba1t(II), nickel(II), and zinc(II) and X is chloride or bromide. The structure of this complex ion is not known exactly; however, it is postulated to be tetrahedral or distorted tetrahedral. 3. Complexes of the type LZMXZ’ where L is either a monodentate or a bidentate ligand, have been reported for a number of metal salts (22, 27-31, 33-41). These complexes have several different structural configurations. The most common structures with the monodentate ligand are the tetrahedral and the polymeric octa- hedral structures with halogen bridges. These structures occur most frequently with pyridine and substituted pyridines (27-31, 33-36), however, complexes with quinoline (30, 31, 35), picoline (30,31), aniline (37b), 2:, ET’ and prtoludines and 2,5-, 2,6- and.3,4-xylidines (37a, 30) have also been reported. The type of structure can be identified by the location of the far-infrared M-X vibrational frequencies. Another possible structure with the same stoichiometry is the §i§_or trans square planar form where the divalent metal ion is platinum(II) or palladium(II) and the organic ligand is pyridine (38-40). A third form, a dimeric (L4M2+)(MX42'), has also been 7 postulated (22); however, no evidence has been found as yet for its occurrence in these heterocyclic systems. 4. A limited number of complexes of the type LMX2 have also been reported. The heterocyclic ligands, such as 1,2-, 1,3-, and 1,4-diazine (42), 1,2,4-triazole (42-44), 1,10-phenanthroline (42) and 2,2'-bipyridine (41, 42) have two donor sites. In most cases the structures are octahedral with halogen and/or ligand bridging. The complexes of the last two types, L MX and LMXZ, as a rule, form 2 2 metal-halide chains as shown below which produce some unusual chemical C1 C1 C1 C1 \M/\/\M/\ M’/// M M \ / \c1/ \m/ \c/ C1 properties. In some cases this chain lowers the magnetic moment of the metal complex much below that found for similar monomeric.structures, while in others the moment remains normal (44). The complexes with the metal-halide chains also retain many of the characteristic prOperties of polymers such as low solubility, high melting points, high molecular weights, and characteristic M-X-M bridged stretching and bending fre- quencies in the far-infrared. II. EXPERIMENTAL PART A. Reagents Pentamethylenetetrazole (PMT) Pentamethylenetetrazole was obtained from Knoll Pharmaceutical Co., Orange, New Jersey, under the registered name of 'Metrazol' and used without further purification. 2,2F-Dimethoxypropane (DMP) Technical grade 2,2'-dimethoxypropane (DMP), obtained from The Dow Chemical Co., Midland, Michigan, was purified by fractional dis- tillation at 28° under 100 mm of pressure. Methanol Reagent grade methanol was refluxed over granulated barium oxide and then distilled at atmospheric pressure. A vapor phase chromatogram of the purified methanol gave a single peak using a Beckman GC-2 chromato- graph equipped with a silicon 30 column (current 200 ma., attenuation 10, temperature 70°). Nitromethane Nitromethane of c.p. grade was passed through a cation-exchange column prepared by activating 75 gm of amberlite IR-120 (acid form) resin in 75 m1 of 0.1 N_hydrochloric acid followed by washing with three portions of 75 m1 anhydrous methanol. Another 75 m1 portion of anhydrous methanol was used to slurry the resin and pack a column 2 cm by 25 cm. The column was then washed with about 500 ml of anhydrous 9 methanol, at a rate of 2 ml per minute, followed by 500 ml of nitromethane which was discarded. The nitromethane to be purified was passed through the column at a rate of 1 to 1.5 ml per minute, then refluxed for 12 hrs. over barium oxide and fractionally distilled through a one meter column with the middle portion being distilled directly into storage bottles. The boiling point of the purified nitromethane was determined to be 101.2° at 760 mm pressure. Other Solvents Reagent grade acetone was dried over Drierite (CaSO4). Methylene chloride, carbon tetrachloride, chloroform, and anhydrous diethylether were used without further purification. Transition Metal Salts The hydrated transition metal chlorides and bromides were obtained from Alfa Inorganic, Inc., Mallinckrodt, and Fisher Scientific Company and were used without further purification. Barium Oxide Barium oxide, used as a drying agent, was obtained from Barium and Strontium Chemicals, Inc., and from Fisher Scientific Company. B. Analytical Methods Elementary Analysis Analysis for carbon, hydrogen, nitrogen, chloride, and bromide was carried out by Spang Microanalytical Laboratory of Ann Arbor, Michigan. 10 Nickel Determination Nickel was determined by standard EDTA titration using Eriochrom Black T indicator. A 300 mg. sample of the complex was transferred to a 100 m1 volumetric flask and dissolved in 75 ml of distilled water before dilution to the calibration mark. A 20 ml aliquot was then titrated by adding 25 m1 of EDTA solution, two or three dr0ps of indicator, and 10 ml of pH 10 buffer followed by back titration with a standard zinc sulfate solution until the color of the solution changed from blue to red. C. Instrumentation Ultraviolet and Visible Spectra The Cary 14 recording spectrometer was used to obtain spectra in the visible and ultraviolet regions. Reflectance Spectra The reflectance spectra of the solid complexes were obtained with a Beckman DK—2 spectrometer with a reflectance attachment at The Dow Chemical Company, Midland, Michigan. Infrared Spectra The infrared spectra in the 5000-650 cm'1 region were obtained with the Unicam 200 spectrometer using Nujol mulls prepared by adding approxi- mately 25-30 mg of ground sample to 10 drops of Nujol. The mulled samples were transferred to sodium chloride plates with a 0.1 mm polyethylene spacer. In some cases samples were also run on the Beckman IR-S spectrometer for better location of band maxima in the 1200-650 cm-1 11 region. Potassium bromide pellets were not used because many of the samples decomposed under the pressure required to make the pellets. The far-infrared Spectra (650-110 cm'l) were obtained with a Perkin- Elmer 301 spectrometer using Nujol mulls. The mulls were prepared as previously described and were pressed between 0.5 or 2 mm polyethylene windows. Magnetic Susceptibilities An Alpha Scientific Laboratories AL 7500 M Electromagnet and AL 7500 P.S. Power Supply equipped with a Mettler single pan balance were used to obtain the magnetic susceptibility measurements. X-ray_Diffraction Photographs A North American Philips Company, type 12045, X-ray generator equipped with a North American Philips Company, type 52026, camera was used to obtain the X-ray powder diffraction photographs. Melting Points A Fisher-Johns Melting Point Apparatus previously calibrated with Arthur H. Thomas Company (Philadelphia, Pa.) micro-melting point standards was used to obtain all melting points. These melting and decomposition points are listed in Table I. D. Preparation of Coordination Compounds of Pentamethylenetetrazole Since PMT is a relatively weak electron donor , water molecules, when present, compete with it for the coordination sites around a metal ion. In fact, PMT complexes of a number of transition metals cannot be prepared 12 in aqueous solution (21a). In order to avoid interferring reactions of water, hydrated metal salts used in this investigation were treated with a dehydrating agent, 2,2'-dimethoxypr0pane, hereafter abbreviated as DMP. Dichlorobis(pentamethylenetetrazole)chromium(II) In order to obtain the complex, chromium(II) chloride had to be pre- pared using a modification of the method outlined in Inorganic Synthesis (45). Preparation of Chromium(II) Acetate Chromium(II) acetate was prepared in a reaction vessel, as shown in Figure l” which was flushed continuously with nitrogen. The sodium acetate solution prepared by dissolving 84 grams of sodium acetate in 100 ml of deoxygenated water was placed in the side arm drOpping funnel, then the entire system was flushed with nitrogen for 15 minutes before the sodium acetate solution was delivered to the erlenmeyer flask. Chromium(III) chloride solution prepared by dissolving 30 grams of chromium(III) chloride hexa- hydrate in 40 m1 of deoxygenated distilled water and 10 ml of 2 N_hydro- chloric acid was passed through a zinc reduction column into the erlenmeyer. The rate of flow was adjusted so as to insure a bright blue chromium(II) solution dropping into the sodium acetate solution. When all of the chromium(III) chloride solution-had been delivered to the erlenmeyer, the suspension of the solid chromium(II) acetate was stirred for 30 minutes and then transferred to a glove box withnitrogen atmosphere. The chromium(II) acetate was filtered under suction and washed with 50 ml of chilled deoxygenated water. Preparation of Chromium(II) Chloride The chromium(II) acetate prepared above was transferred to a chilled Figure 1. 0 v ------- (d) (‘L’T‘(’('1’(’Tr(’('( Reaction Vessel for Chromiumflll Acetgtg a) b) e) d) 500 ml erlenmeyer flask y-connection with ground glass fittings 500 ml drapping funnel with pressure equalizing side arm 50 ml buret filled with 20 mesh zinc granules 14 250 ml erlenmeyer flask containing 60 ml of concentrated hydrochloric acid which was held at 0° by a salt-ice bath. A light blue solid chromium(II) chloride tetrahydrate was formed. The supernatant liquid was decanted and the chromium(II) chloride was washed with three 25 ml portions of chilled concentrated hydrochloric acid. Preparation of PMT Complex The crystalline.chromium(II) chloride tetrahydrate was added to 75 ml of a 50-50 UMP-methanol solvent mixture to remove the waters of hydration. Assuming 100% yield of chromium(II) chloride a 6:1 ratio of PMT to salt was added to the chromium(II) chloride solution and allowed to stir until the solid blue complex was formed. The complex was fil- tered, washed with DMP, dried under vacuum, and stored in nitrogen atmos- phere. The stability of the complex was tested by exposing a sample to moist atmosphere. The sample rapidly changed color from blue to green indicating formation of chromium(III) chloride. The oxidation also occurred over a period of several days, in a dry atmosphere. Analysis and physical prOperties for the above and all succeeding complexes are listed in Table I. Dichloro(pentamethylenetetrazole)manganese(II) The hydrated transition metal chloride (0.02 mole) was added to 75 m1 of 50% UMP—methanol solvent mixture and stirred for 10-15 minutes to remove the waters of hydration. A slight excess of PMT (0.04 mole) was dissolved in 25 ml of the same solvent and added slowly from a dropping funnel. A finely divided pale pink. powder was formed which was filtered, washed with DMP, and dried under vacuum. 15 Dichloro(pentamethylenetetrazole)iron(II),-cobalt(II), -nickel(II), and -copper(II) These complexes were prepared by exactly the same procedure as the dichloro(PMT)manganese(II) complex. Finely divided yellow, pink, light green, and light green powders were obtained respectively for iron(II) cobalt(II), nickel(II), and c0pper(II). The latter two complexes were also prepared by dissolving nickel chloride and c0pper chloride hexahydrate (0.02 mole),respectively, in minimum amounts of water and adding slowly to a solution of PMT (0.03 mole) in 75 ml of glacial acetic acid. Three ml of acetic anhydride was added and the solution was stirred for one-half hour. A light green powder formed in each case. To remove uncoordinated PMT a Soxhlet extractor with carbon tetrachloride was used to purify the complexes. The decomposition points of these complexes were checked with those pre- pared previously and_agreed very well. A comparison of the yield for the two methods in the case of the copper system showed it to be 97.5% for the DMP-methanol method and 93.1% for the acetic acid method. Dichlorobis(pentamethylenetetrazole)zinc(II) The procedure outlined for the dichloro(PMT)manganese(II) complex was followed. However, the product was a viscous oil. The oil was dis- solved in dichloromethane and the solution was added slowly to anhydrous diethylether. 1\fine1y divided white precipitation was formed. In order to obtain pure product the precipitation procedure was repeated several times. 16 Dibromo(pentamethylenetetrazole)manganese(II) The hydrated manganese(II) bromide (0.02 mole) was dissolved in 75 m1 of 50% DMP-methanol solvent mixture and stirred for 10-15 minutes to remove the water of hydration. A slight excess of PMT (0.04 mole) dissolved in 25 ml of the same solvent mixture was added slowly from a dropping funnel. A finely divided pale pink complex was obtained which was filtered, washed with DMP, and dried under vacuum. Dibromobis(pentamethylenetetrazole)cobalt(II) and -zinc(II) These complexes were prepared exactly as the dibromo(PMT)manganese(II) complex. Crystalline royal blue and white powders were obtained respectively for cobalt(II) and zinc(II) complexes. Dibromo(pentamethylenetetrazo1e)nickel(II) Since anhydrous nickel bromide was not very soluble in the solvent mixture used in previous syntheses, the salt (0.02 mole) was dissolved in 75 ml of 90% acetone-water mixture. A slight excess of PMT (0.04 mole), also dissolved in acetone, was added slowly from a drOpping funnel. The system changed color but no solid formed until the volume was reduced and the solution added to chloroform. A light green solid complex was isolated. Dibromo(pentamethylenetetrazole)copper(II) This complex was prepared exactly like the dibromo(PMT)manganese(II) complex. A light rust brown complex was obtained. Pentamethylenetetrazole Complexes with Chromium(lll) Chloride Iron(III) Chloride, and Iron(II) Bromide Attempts were made to prepare these complexes by the techniques 17 described above, but they were unsuccessful. In the case of the chromium(III) and iron(III) chlorides no evidence for complex formation could be obtained. The reaction of iron(II) bromide with PMT initially yielded an oil which dissolved in excess DMP. When this DMP solution was added to anhydrous diethylether a solid was isolated; however, analysis showed that the compound was non-stoichiometric. In a number of the above described preparations, the PMT to transi- tion metal halide ratio was varied in the reaction mixture and in all cases the stoichiometry of the resulting complexes was not affected. E. Experimental Studies Solubility and Thermostability_ Complexes having one PMT molecule per metal ion were found to be quite insoluble in benzene, nitromethane, chloroform, carbon tetrachloride, formic acid, and acetic acid. They were somewhat soluble in acetoe nitrile, methanol, ethanol, dimethylformamide, and pyridine. Spectro- photometric evidence, however, indicates that the dissolution is accompanied by the break-down of the complexes. This behavior is not unexpected in view of the relatively strong donor pr0perties of these solvents and their ability to compete with PMT for the coordination sites around the metal ion. In the cases of the PMT complexes with cobalt chloride and bromide, as well as copper chloride and bromide in pyridine, the solid complexes changed color but remained insoluble, indicating possible ligand replacement at the surface of the solids. The complexes with two PMT molecules per metal ion were insoluble in benzene, chloroform, and carbon tetrachloride and slightly soluble in l8 nw.Hm nn.am mm.~m mH.- No.e mo.e mn.wm oo.w~ mafia: Hm.vv m~.mv om.ma eo.mH mn.~ en.m mm.m~ ow.mH cache pmsm Hw.ee mo.ve H5.ma oc.m~ mw.m oo.m om.o~ mm.mH :oonm :oHHe> mm.mm mm.~m mo.~m no.- no.e no.e 5H.mm oo.mm mafia Hexom m~.me Hm.me mw.ma me.mH om.m wn.~ Ne.om oe.o~ xcfim .uq wH.uH v~.nH o~.n~ N~.nm ww.e mm.e mm.em mm.em mums: Ho.om mm.mm om.o~ ne.o~ on.m en.m mv.om wm.om coenw rug me.o~ oe.o~ mm.o~ Nw.om ou.m ew.m Hm.om mo.s~ coeuw .94 oe.om oe.o~ Hm.om ow.o~ on.m on.m mm.om em.o~ gang 55.0N mu.om mH.HN NN.HN ow.m vw.m om.n~ wH.uN 30~Ho> ow.om on.o~ mm.HN mo.H~ Nw.m mw.m mm.nm nm.nm Mafia mama on.n~ cw.nH no.wm mo.m~ mo.m mm.e oH.om eo.om :eenm asap .pq .oamu \mcsom .oHeu .wcaom .uaao .eenom .oHeu venom noaou x» 2* aw ow momxaec< oHoNeHueuesoneueamucom guwz mexenmaou vegan: Heuoz :oflpwmaehh mo mefiuHQQOHm meow ecu momxamc< oo-emH oomH.v ooomA oHoHummH ooomA cannoH oomN.v ooomA oOCMA oovH.v ooomx oooH.v .4mqm .H Queue NummfiazaueN NHmHAAZavsu NumHAHZQUnz mammfiezavou Numumezavcz «HomfiazaveN NHanazavzu Nuowezavnz Nuuuflezavou Nuofiflezavoa muufifiezavez NHUNAAZQUHU xoflmsoo 19 nitromethane and acetonitrile. The two classes of complexes strongly differ in their decomposition temperatures. Those with the structure MII(PMT)1X2 in general are quite stable thermally and decompose at 100-150° higher than the MII(PMT)2X2 compounds, as shown in Table I. Mole-Ratio Studies and Formation Constant Determination The method of mole-ratio is used for the SpectrOphotometric deter- mination of the stoichiometry of a complex in solution, especially when only one complex is formed. The total analytical concentration of metal (or ligand) in a series of solutions is held constant while the concentra- tion of the other component is progressively increased. Absorption spectrum of each solution is taken and a plot of absorbance 12, mole ratio of ligand to metal concentration is obtained. Under favorable circumstances (gpgp, when only one complex absorbs at a given wavelength), it is possible to determine the ratio of ligand to metal in the complex. This method was applied to the c6balt(II) chloride-PMT system using purified nitromethane as the solvent. The solutions to be studied had a constant cdbalt(II) chloride concentration of 2 X 10'4 M_prepared by using 10 ml of 5 X 10'4 M c6balt(II) chloride stock solution and a varied amount from 0.2 to 6.0 ml of 5 X 10'3 M_PMT stock solution followed by dilution to 25 ml with nitromethane. At concentrations greater than 1:1 mole ratio the solutions became cloudy and solid complexes were precipitated. Because of the limited solubility of the above and the other PMT complexes, the stiochiometry and formation constants could not be obtained spectrOphotometrically. However, the fact that 20 precipitation occurs at a given mole ratio (e.g. 1:1 for the cobalt(II) chloride-PMT system) indicates that the stoichiometry of the complex has been attained. Composition Study in Nicke1(II) Chloride Systems The PMT to nickel(II) chloride ratio was varied in the DMP-methanol reaction mixture from 1:1 to 6:1 by varying the amount of PMT (0.01 to 0.06 mole) and keeping the nickel(II) chloride concentration constant (0.01 mole). Each complex formed was filtered, washed with DMP, and dried under vacuum. To determine if each complex had the same composition the percent of nickel(II) was determined by titration with standard EDTA (p. 10) and compared to the percent of nickel(II) calculated for pure nickel(II) chloride (45.30%), dichloro(PMT)nickel(II) (21.93%), and dichlorobis(PMT)nickel(II) (14.46%). The data obtained are shown in Table II. The percent of nickel ranges from 18.52 to 22.67 with an average value of 20.24 which is most indicative of the composition dichloro(PMT)nickel(II) rather than any other form. Reflectance Spectra Due to the insolubility of the complexes, it was impossible to study the solution Spectra, thus reflectance Spectra of the powders were obtained using magnesium carbonate as the reference. The observed absorption band maxima are listed in Table III along with literature data on the principal absorption bands of the corresponding octahedral hexaaquo complex ions and the respective tetrahalide complexes (tetrahedral or square planar configuration) for comparison purposes. Table II. Composition Study of PMT-Nickel(II) Chloride Systems Sample Ratio Sample Size % Nickel (PMT:NiC12) (Grams) (Found) 1 : 1 0.3007 19.22 2 : 1 0.2987 18.55 3 : 1 0.3161 18.52 4 : 1 0.3102 22.24 6 : 1 0.3030 22.67 Average 20.24 [(mlEDTA) (NEDTA) '(mIZnSO4) (NZnSO4)] (meg. wt. N1) X 100 % N1 = (Sample weight) (0.20) NanO = 0.0102 EDTA 0.0096 22 X-Ray Diffraction Photographs Attempts were made to obtain X-ray diffraction phOtographs of the powders. However, the patterns obtained were.very poor, consisting of numerous weak lines; thus, the method was abandoned. Magnetic Susceptibilipy The magnetic susceptibilities of the transition metal complexes were obtained by the Gouy method.. This method.is based on a comparison of the weight of a substance suspended so that half of it is in a strong uni- form magnetic field to the weight of the.sUbstance free.of a magnetic field (Aw). This weight change can be related to the magnetic suscepti- bility of the complex and to the.number of unpaired electrons, n, through a series of equations as.described below. The.gram susceptibility, xg, the induced moment per gram per unit field, for.any substance is cal- culated from the following relationship and has units of cm:5 gm‘l. 2 A'w = ASH ) = C (1.) 9 X3 2 ' A (H2) = C = apparatus constant in grams 2 A = Cross sectional area of Gouy tube in cm2 H = Strength of magnetic field in gauss Aw = Change in weight of substance in and out of the magnetic field in grams 9 = Density of the substance in gm cm'3 = Gram susceptibility of the substance in cm3 gm"1 Complex Mn(PMT)1C12 Mn(PMT)lBr2 Fe(PMT)1ClZ Co(PMT)1C12 Co(PMT)zBr2 Table III. Observed Bands £21 355 376 426 472 560 1,190 1,390 360 379 432 473 554 1,200 1,390 1,080 1,412 508 560 692 1,500 387 432 447 524 628 1,095 1,395 -1 cm 28,200 26,600 23,500 21,200 17,900 8,200 7,200 27,800 26,400 23,100 21,100 18,000 8,300 7,200 9,300 7,100 19,700 17,900 14,400 6,700 25,800 23,100 22,400 19,100 15,900 9,100 7,200 (VS) 23 Reflectance Spectra Lit. Values (solution Spectra) Octahedral* 29,700 28,000 25,150Ca) 24,900 23,000 18,800 10,400(°) 21,550(b) 19,400 16,000 8,000 Tetrahedral or Planar** 23,250(33 22,500 27,000(d) 23,000 22,100 21,500 Charge transfer beyond 20,000 (g) 16,900(d) 16,300 15,700 15,000 14,400 16,200(d) 15,500 15,000 14,300 13,800 24 Table III (continued) Observed Bands Lit. Values (solution spectra) Complex mp. cm"1 Octahedral* Tetrahedral or Planar** Ni(PMT)1Cl2 422 23,700 25,500(a) 16,400(d) 698 14,300 18,500 15,150 1,190 8,400 15,400 14,300 1,435 7,000 13,500 1,935 5,200 8,600 Ni(PMT)lBr2 435 23,000 16,500(f) 720 13,900 15,300 1,195 8,400 14,300 1,455 6,900 1,945 5,100 Cu(PMT)1012 780 12,800 12,600(°) 27,800(d) 23,000 11,200 Cu(PMT)lBr2 800 12,500 27,100(d) 23,100 15,600 11,500 a: II 2+ Spectra for octahedral aquo complexes M (H20)6 . ** Spectra for the corresponding tetrahalides MII 42'. (a) T. M. Dunn, in'Modern Coordination Chemistry',‘ J. Lewis and R. G. Wilkins, Editors, Interscience, New York (1960), p. 229 ff. (b) C. K. Jdrgensen, Acta. Chim. Scand., 9, 116 (1955). (c) C. K. Jdrgensen, ibid., S, 1502 (1954). (d) R. s. Nyholm and N. s. 0111, J. Chem. SOC., 3997 (1959). (e) N. S. Gill, ibid., 3512 (1961). 25 The gram susceptibility is related to the volume.susceptibility, X, a dimensionless quantity, through the density of the substance. This vol- ume susceptibility is the induced moment per unit volume per applied field. X = Xl/p (2.) The molar susceptibility (Xm), the induced moment per mole per applied field, has dimensions of cm3 mole'1 and can be expressed as the product of the gram susceptibility and molecular weight (M.W.). x, (M.W.) (3.) p x = ‘xg (M.W.) The molar susceptibility is also related to the magnetic moment by the following relationship xm = N 82112 (4.) 3 E T where: = Avogadro's number = 0.917 x 10’20 Magnetic moment erg gauss"1 Bohr magneton = Boltzmann's constant F-IW'CWZ II = Absolute temperature Rearrangement of equation (4) and substitution of numerical values for the constants gives the magnetic moment in Bohr magnetons. 3 k xm'r 1/2 “ - [ ] = 2.84 (E. Tf/2(5.) N 82 Assuming spin-only approximations the number of unpaired electrons (n), 26 can be related to the magnetic moment by the following relationship 1/2 u = [n(n+2)] (6.) The Gouy susceptibility tube used in this study was constructed from 6 mm I.D. Pyrex tubing. The distance from the septum to the reference mark was 60 mm. The volume of the tube (2.47821 cms) was calculated by measuring the weight of distilled water (2.47181 gm) needed to fill the tube to the reference mark and dividing by the den- sity of water (0.997418 gm cm's) at the ambient temperature of the water (23.5°). The Gouy apparatus was calibrated using two reference materials, mercury(II) tetrathiocyanatocobaltate(II) and copper(II) sulfate pentahydrate. These reference materials, as well as the finely ground PMT complexes, were dried and packed into the Gouy susceptibility tube by placing a small quantity of the sample in the tube and tapping it firmly on a stone surface sixty times. This process was repeated until the tube was filled to the reference mark. The tappings were used to minimize the error resulting from lack of uniform packing, which is the main source of error in the measurement. The Gouy susceptibility tube was attached by a copper wire to the pan of a Mettler single pan balance and sus- pended between the pole faces of an Alpha Scientific Laboratories electromagnet. All changes in weight were recorded at three field strengths with the average of three successive measurements being taken to prepare a calibration graph. From the calibration graph of weight change 23: % field strength, the value of weight change at 50% field strength was used in all the calculations. Those values are listed in 27 Table IV along with the other data. Mercury(II) tetrathiocyanatocobaltate(II) and cepper(II) su1fate pentahydrate are reported by Figgis and Nyholm (47) to have gram sus- 6 6 cm3 gm.1 respectively ceptibilities of 16.44 (i 0.5%) X 10- and 5.92 X 10' at 20° and to obey the Curie-Weiss law between 10-30°. The copper(II) sulfate pentahydrate is known to be difficult to pack uniformly,which leads to some error in reproducibility. Since all the measurements were obtained at room temperature of 23.5° the gram susceptibilities of the reference materials were cor- rected as follows: -6 3 -1 296.65°K X _______. pg Hg[Co(SCN)4] 16.44 X 10 cm. gm 293.1s° K = 16.64 x 10‘6 Cm? gm—l O ’% CuSO .5 H 0 = 5.92 x 10'6 cm3 gm-l 296.65 K 4 2 -6 3 _1 293.150 K = 5.991 x 10 cm, gm From the data in Table IV the apparatus constant (C) in equation (1) was calculated for both reference materials. . A " C ref. = ref. 9 ref. Xg ref. 0.1080gm Hg[0°(5CN341 1.219 gm cm3 16.64 x 10‘6 cm3 gm‘1 = 5,346.9 gm 0 CuSO4-5 H20 = pp 3 0.0412 _6 3 -1 1.2786 gm cm 5.991 x 10 cm gm = 5,373.1 gm 28 Combining equations (1) and (3) the observed molar susceptibility of each of the complexes was calculated using the data in Table IA! and the above calculated apparatus constants. X = X = A m obs. ‘ g (M.W.) w (M.W.) C o In order 11) obtain the true value of the paramagnetic contribu- tion to the observed molar susceptibility, a diamagnetic correction must be applied to account for the paired electrons in the ligand, halide ion, and central metal ion. The respective molar diamagnetic suscepti- bility corrections are listed in Table \l and are applied in the following manner: X - X b X _ X X - m corr. m obs.+ a m (M2+) + m (PMT) + c m (X ) where a, b, and c are the re5pective subscripts of the empirical formula II of the complex M a(PMT)ch. Infrared Spectra of Pentamethyleneteprazole and Its Complexes with Transition Metal Halides The infrared spectra of' PMT and the PMT complexes with transi- tion metal halides were obtained in the 5000—650 cm'1 region. The spectra of these complexes were very similar. The variation in the frequency of the band maxima was about *5 cm'1 depending on central metal ion, halide, and number of PMT molecules per metal ion. Figures 1-5 (in Appendix I) are representative of uncomplexed PMT and the respect- ive chloro and bromo complexes containing one and two PMT molecules per metal ion. A comparison of the Spectrum of uncomplexed PMT with the 29 nemme.e mwsmm-e £56A5.8 «mama 5 naesm.m auawm.w emnum.oa ameem.oH emwmu.mu meeeN.mH nommm.ms maawm.mu psoom.uu amquw.uu nummm.w mamee.w Nmom.v MI OH x fl.ue66V5x Mowmw.e Ao.emm nmuoa.e moema.e AA.A6N amnmo.w mmmme.w Hu.mme emamn.en mNNmN.HH No wow pmeom.mu «swam.mu mm.mmm nemHe.~H asmae.mn no.56~ nmoem.NH «AHON.NH mo.eo~ nfiwmw.m enumm.w em.mmm E m-oH.x_x .32 .Hoz mums xwfiHflnwumoomsm owuecwmz mmoo.o emvo.o omwo.o mmoH.o owmo.o oomH.o veno.o oovo.o Navo.o oon.o 3 < .>H manme NwNmH.N eonnH.H nnmmm.m muowH.H NHHwH.H wmemN.H onmmn.o vnwmw.o ammofi.m enmoo.m .uz mameew NemufiezaVez Nfiufifiazavuz NemNAAZAVou NHUHAPZAVOU NemHfiEficz NHuHmazavcz Nfiuflhazavma NHUNAazaVeu ON: m.eomeu Heflzomvoung mfimsem 30 00.0 e . 000 0 n00.0 000 0 nem.u amm.H nmm 0 0 H «mm H N v o z m. omSU Hmflhouez euceeemmm fienzomvou_m: Hefluopmz eucoeomem --- --- mm.H0m --- --- 00.0mm nanmm.o ammom.H 00000.H mm0fim.s 00.nem 0m000.0 e000N.H a0000; 00N0~.H ~0.NAN E TS x 7880 Ex 0 2 x x .03 .3: mmoo.o: omoo.o| mmoo.o NvHo.o II .D ll (6 mummm.m omnnw.a vaHN.H mewm.H .pz onsem N00NA02000N Naommaza0cN N H 00 mazavsu NEH £2008 eHmEmm “0.30660 >H oases 31 Table V. Molar Susceptibility Corrections for Diamagnetic Behavior Ion xm(corr.) X 10'6 * Crz+ 15.0 Mn2+ 14.0 Co2+ 12.8 N12+ 12.8 Cu2+ 12.8 Zn2+ 15.0 Fe2+ 12.8 01‘ 26.0 Br- 36.0 PMT 222.0 *Reference (48) 32 spectrum of a representative complex, reveals two noticeable changes in the spectrum of PMT upon coordination. The quartet of bands at 1100 cm-1 was split into two sets of bands,and the band at 1000 cm.1 disappeared. A literature review of previously reported luff complexes indicated that the same Split in the quartet of bands at 1100 cm-1 was observed by Holm in the bis(PMT)silver(I) nitrateecomplexes (19); however, in the first row transition metal perchlorate complexes studied by D'Itri (21a) the free perchlorate ion has a broad band in this area which covers any changes in the spectrum of PMT. Theafar-infrared Spectra of PMT and its complexes were obtained in the 650-110 cm'1 region. The Spectrum of PMT shows a total of nine absorption bands in the 700-180 cm'1 region as reported previously by D'Itri (21) with an additional broad band at £23 150 cm-1. The relative intensities of the bands observed and the location of the absorption maxima are represented by bar graphs in Figures 2 and 3. A comparison of the spectra of the complexes with the Spectrum of uncomplexed PMT shows that there are four bands between 700-425 cm”1 which are apparently ligand vibrations. In this region the dichlorobis(PMT)chromium(II) and the dibromo(PMT)c0pper(II) complexes have a weak band at about 560 cm-1. Below 425 cm'1 the complexes all exhibit numerous bands, some of which are no doubt ligand vibrations. The other bands are probably due to metal-halogen and metal-nitrogen vibrations upon complexation as well as lattice modes. Because of the complexity of the spectra it is impossible to make any unambiguous band assignments. Some very tenta- tive generalizations, however, can be obtained by comparing these spectra with those of similar complexes. 33 .mfifisz Hohsz .fifiuao :09 mexodmaou opwHoHnu mo ehuuemm nowumuomn< veaehmeHuumm .N enamHm '34 OOH 00m 00m 000 000 000 00A 0 _ p — . . p _ . _ ESQ - _ _ _ _ _ _ _ _ _ - _ __ : _ _ _ _ - _ +02 _ .1: _ ,0 _ 0 _ _ _ L . +NOU _ __ w : n _ _ _ _ _ +N0n— C : ,, _ _ _ _. _ _ _ _. _ - . as +HU w _ _ _ - _ _ _ _ N _ — . — e _ — _ _ - _ 35 .312 Honsz .nfiuao :3 moxeflmsou 035on we «300mm 539334. woneamcHuymm .m chem: >36 ooH com com oov com 000 com +N¢N __H~ _ _.__ _ +Nsu -_ _. L ___ _ _ : _ . _ _ x .N. A _ 1: _ _ ._ :2 _ ._ __ __P _ III. DISCUSSION AND RESULTS The replacement of the perchlorate salt of the transition metal by a halide resulted in a significant change in the stoichiometry and the structure of the PMT complexes. Unlike the perchlorate complexes which had six PMT molecules per metal ion, the chloro and bromo com- plexes have only one or two PMT molecules per metal ion, forming two new types of PMT complexes. The first type of complex contains two PMT molecules per metal ien.and consists of the chloro complexes of chromium(II) and zinc(II) and the bromo complexes of cobalt(II) and zinc(II). These complexes melt or decompose at much lower temperatures than those with one PMT molecule per metal ion and tend to be more soluble in some solvents. In the case of the cobalt(II) complex the measured magnetic moment (4.48 B.M.) was in the range indicated by Figgis and Lewis (49) for a cobalt(II) ion in a tetrahedral environment (4.30-4.80 B.M.) rather than in an octahedral environment (4.80-5.20 B.M.). Likewise, the reflectance spectrum of the complex indicated that the cobalt(II) ion is in a tetrahedral environment with very intense absorption bands at 15,900, 9,100, and 7,200 cm'1 which are similar to the intense bands observed by Nyholm and Gill (46d) for the tetrahedral tetrabromo- cobaltate(lI) ion. Since the zinc(II) complexes are white in color and do not have unpaired electrons, very little information could be obtained regard- ing the environment of the central metal ion. It must be noted, however, that zinc(II) ion has no ligand field stabilization effects because of 37 38 the completed d shell, thus the stereochemistry of its complexes has to be determined solely by considering the size and the bonding forces of the ligands. Zinc(II) is known to form complexes having coordination numbers of two, four, and six. The coordination number of four is the most common one with the tetrahedral configuration predominating over the square planar. Complexes with coordination numbers of six, such as the hexakis(PMT)zinc(II) perchlorate (21), hexaaminezinc(II) and hexaaquozinc(II) (50), and of two, such as dimethylzinc(II) (50) have also been prepared. Most complexes of chromium(II) have a coordination number of six, such as CrX -2 L, where X is chloride or iodide and L is 2 hydrazine or dipridyl (51); however, there are some exceptions, such as salts, like [Cr (CO)2 diarSZX]X, where X is a halide and diars repre- sents p-phenylenebisdimethylarsine (52), and addition compounds, CrXZ-n NH3 (n I 6, 5, 3, 2, l) where X is a halide (51). At this time it is rather difficult to state with any certainty the exact structure of the zinc(II) and chromium(II) complexes with PMT. However, by analogy to the stoichiometry and physical prOperties of the dibromobis(PMT)cobalt(II) complex, a prediction that the zinc(II) and chromium(II) ions are in either a tetrahedral environment or else in an addition type compound with a structure of MX2-2 PMT, seems reasonable for the dichlorobis(PMT) zinc(II), dibromobis(PMT)zinc(II), and dichlorobis(PMT)chromium(II) complexes. The second type of PMT complexes contains one PMT molecule per metal ion. This type consists of the chloro complexes of manganese(II), iron(II), cobalt(II), nickel(II), and copper(II) and the bromo complexes of manganese(II), nickel(II), and copper(II). These complexes exhibit high melting or decomposition points and are insoluble in most solvents. 39 To obtain some insight as to the environment of the central metal ion of these complexes, the reflectance spectra were obtained because the limited solubility hindered solution studies. The location of the observed reflectance spectra bands are compared in Table III to the corresponding hexaaquo and tetrahalide divalent metal complexes. It Should be noted that in the case of manganese(II) complexes the ob- served bands are rather numerous and cover the same spectral range from 30,000 to 18,000 cm"1 as reported by Dunn (46a) for the hexaaquo- manganese(ll) complex, thus indicating that the central metal ion is. in an octahedral environment. Likewise for the nickel(II) complexes, there are no characteristic tetrahedral bands between 16,500 and 14,000 cm”1 as reported by Nyholm and Gill (46d) for the tetrahalo- nickelate(II) ion, but there are bands at 23,700 and 8,400 cm"1 in the bromo complex which correspond to those reported by Dunn (46a) for the octahedral complex of hexaaquonickel(II) ion with bands at 25,500, 13,500, and 8,600 cm’l. The copper(II) complexes also exhibit 1 one band at 12,800 cm- for the chloro complex and one band at 12,500 1 band cm'1 for the bromo complex which correspond to the 12,600 cm- reported by Dunn (46a) for the octahedral environment of copper(II) in the hexaaquocOpper(II) ion. In the dichloro(PMT)iron(II) complex it is quite evident that the band at 9,300 cm"1 corresponds better to the 10,400 cm'1 band observed by Jdrgensen (46c) in the hexaaquoiron(II) complex rather than to the charge transfer bands beyond 20,000 cm“1 as reported by Gill (46e) for the tetrachloroiron(II) ion. The dichloro- (PMT)cobalt(II) complex, on the other hand, has very weak bands with more Similarity to the absorption bands observed by Jdrgensen (46b) for 40 the hexaaquocobalt(ll) ion than the intense bands observed by Nyholm and Gill (46d) for the tetrachlorocobaltate(II) ion. From the comparison of the absorption bands, it was seen that the central metal ion in all cases appears to be in an octahedral or distorted octahedral environment. Magnetic susceptibility data.indicate- that the complexes were all of the high spin type- In.the.case.of.cobalt((l) where the configuration and magnetic moment range has been.specified by Figgis and.Lewis (49) the moment observed for.the dichloro(PMT)cobalt(II)-of-5.12 B.M. falls within the range specified for octahedral.configuration.of-4.80-5.20 B.M. Both of the dichloro and dibromo(PMT)copper(II) complexes exhibited subnormal magnetic moments of 1.55 B.M. Copper(II) compounds with subnormal magnetic moments.(those below 1.70 B.M.) have been.reviewed byKato, e£_el;, (53) and classified.according to copperetoecopper mag- netic interactions into two classes, direct interactions and super- exchange interactions- In the first class the subnormal.moment is attributed to copper-copper magnetic interaction due to direct bonding between the copper(II) ions as in the case of copper(II) acetate monohydrate and its homologs..-Inuthe.second.c1ass the.copper-to-copper magnetic interaction is due to so-called "super:exchange” interactions. Most of the complexes.belonging to this classlcontain monoatomic bridges having larger copper-copper distances than the.direct.copper-to-copper interaction.distances as in Class I. -Amcnglthese complexes are the halogen-bridged.complexes which-maymhave-normal or subnormal.magnetic .moments. Such dimeric structures as those confirmed by x-ray analysis for KCuCl3 by Willett (54) as well as the anhydrous copper(II) chloride 41 and bromide show normal magnetic moments of 1.77, 1.75 and 1.31 B.M. x‘//Cu\\\x’//Cu\\\x’//Cu\\\X’//Cq\\\x \/ i\c./ N/ i\../ I u I C /C'\§(/ u\;/Cu\1x \\\Cu//'A\\th/’ \\\Cu(// (55, 56, 57), respectively, at room temperature, with the copper(II) x-- bromide having the lowest moment. On the other hand, lnoue,et_§ln, (44) have reported magnetic moments ranging from 1.41-1.89 B.M. for dichloro(l,2,4-triazole)c0pper(II) and related c0pper(II) complexes with heterocyclic ligands. Again, these latter complexes are known to have halogen bridging, but the authors attribute the low moments to Spin coupling through the heterocyclic ligand. In our COpper complexes we have no definite evidence for metal- metal bonds but there is the possibility of halogen or heterocyclic ligand bridging. Reviewing the above properties: (1) stoichiometry of MII(PMT)1X2, (2) high melting or decomposition points, (3) insolubility except in donor solvents, (4) central metal ion in octahedral environment, and (5) subnormal magnetic moments in the case of capper without support for metal-metal bonds, it was evident that the complexes were polymeric in nature with possible halogen and/or ligand bridges. Since it is possible to distinguish between polymeric and mono- meric transition metal halide complexes of pyridine on the basis of 42 differences in the vibrational frequencies of metal—halogen and metal- ligand bonds, and since PMT should Show some similarity to pyridine as a ligand, a comparison of the far-infrared Spectra for the two types of complexes could be fruitful. A variety of pyridine complexes with several different configurations are formed in metal halide systems, among these are the tetrahedral, octahedral, distorted polymeric octa- hedral, polymeric octahedral and planar configurations. Clark and Williams (30) have found that for the distorted polymeric octahedral complexes with composition CuX2°2 Py (X = Cl, Br) there are two metal- 1 halogen vibrational modes, Cu-Cl at 294 and 235 cm- while the Cu-Br vibrations occur at 255 and 202 cm-1. In the case of polymeric octa- hedral complexes of MC12°2 Py (M = Mn, Fe, Co, and Ni), the spectra consisted of broad,badly resolved bands in the 260-200 cm-1 spectral region. Goldstein, 3: el,, (35) have reported that in-plane-bending modes for distorted polymeric octahedral complexes of Cqu'Z Py occur 1 at 175 cm' for the chloro complex and at 130 cm-1 for the bromo complex. Ahuja, eghgl3, (37a) who have studied far-infrared spectra of bridged complexes with the general formula LIMX2 and LZMX2 (where M = Mn, Co, Ni, Cu, Zn and Cd; X = halogen; and L = monodentate hetero- 1 cyclic amines) reported M-Cl vibrations occur at £23 230 cm' and M-Br vibrations are located below 200 cm'l. Purely by comparison of metal halide complexes of pyridine (MX2‘2 Py distorted polymeric octahedral) and PMT (Table VI), the bands at 206, 253, 300 and 299 cm-1 can be assigned to the metal-halogen vibrations of the respective Mn(PMT)1C12, Ni(PMT)1C12, Cu(PMT)1Cl2 and Zn(PMT)2Cl2 complexes. In turn, if the bands at 163, 198, 245 and 43 220 cm.1 are assigned to the metal-bromide vibrations of the respective Mn(PMT)1Br2, Ni(PMT)lBr2, Cu(PMT)lBr2 and Zn(PMT)ZBr2 complexes, the ratio of the frequencies of M-Br to M-Cl are 0.79, 0.78, 0.81 and 0.74, respectively, which is Similar to the value of 0.81.calculated on the basis of reduced mass changes and no change in force constant. Inspection of the spectra also reveals that Ni(PMT)1C12 and Ni(PMT)lBr2 complexes both have a band at 174 cm"1 which cannot be ascribed to ligand vibration. This band probably represents metal- ligand vibration. Using similar reasoning, the bands at 220 and 190 cm"1 correspond to metal-ligand vibrations for the Cu(PMT)1X2 and Zn(PMT)2X2 complexes, respectively. This is in line with tentative I made by D'Itri for the metal-ligand vibrations ranging from 236-198 cm- PMT complexes with transition metal perchlorates (21a). In the case of cobalt(II), no comparison between the spectra of the chloro and bromo complexes could be made since they belong to different types of structure. Since the bromo complexes of chromium(II) and iron(II) were not prepared, the data collected were not adequate for interpre- tation of the far-infrared spectra. The manganese complexes, on the other hand, did not provide information which would allow even tenuous metal-nitrogen band assignments. It is evident that the metal halide complexes, containing only one molecule of PMT, even with bridged halogens, do not entirely satisfy any type of structural stereochemistry in which the.metal ion is in an octahedral environment unless there are.metal-metal bonds or bridging ligands. Again, the evidence for metal-metal bonds is, at best, very tenuous 44 Table VI. Metal-Ligand and Metal-Halide Vibrational Frequencies (cm-1) Compound \) (M-L) \’(M-X) \’ (M-Br)/"(M-C1) Mn(PMT)1C12 198 206 0.79 Mn(PMT)iBr2 195 163 Mnc12 2 Pya 212 233 MnBr2-2 Pya 212 <200 Ni(PMT)1C12 174 253 0.78 Ni(PMT)1Br2 174 198 NiC12°2 Pya --- 246 NiBr2°2 Pya --- <200 Cu(PMT)1C12 220 300 0.81 Cu(PMT)1Br2 220 24s CuC12-2 Pya 268 294 CuBrZ-Z Pya 269 255 Zn(PMT)1C12 195 299 0.74 Zn(PMT)lBr2 190 220 a = Reference (30) 45 for the c0pper complexes and nonexistant for the other complexes. On the other hand, in all previous work PMT was found to act as a mono- dentate ligand. The reluctance of the tetrazole ring to form more than one bond is illustrated by the PMT—halogen complexes,such as PMT:IC1, where, as shown by recent crystallographic study(l8), iodine monochloride is bound to N-4 nitrogen. Nevertheless, it seems that in the case of metal halides the polymeric form, due to halogen bridging, forces the tetrazole ring into a bridging position. A similar phenomenon was observed by Jarvis (43) when the crystallographic study of dichloro- (l,2,4-triazole)copper(II) complex was determined as shown in Figure‘l. Here the c0pper(II) ion finds itself in an octahedral or distorted polymeric octahedral environment with bridged chlorine atoms forcing the triazole ring into a bridging position. By analogy to this complex, it is possible that the complexes prepared in this study containing only one molecule of PMT per metal ion represent the first case where PMT acts as a bridging ligand. 46 Figure 4. Dichlorofi,2,4e‘triazoie)copper('ll) Complex jStructural Unit H I / N HC \ CH N N. C1 1 \\~\ ‘,///””’ \"“~ ‘,1,/”””C °\"‘- ,///’//” \ ‘ \ ‘ Cu Cu ‘Cu \ \ \ / \ \ / \ T / \ \ \ \ \\ C1 C1 10. 11. 12. 13. 14. 15. 16. 17. 18. REFERENCES Oliveri-Mandala', E., Gazz. Chim.NItal., 34, 175 (1914). Moore, 0. W. and Wittaker, A. G., g, An, Chem. Soc., S3, 5007 (1960). Lounsbury, J. B., Q, Phy . Chem., 92, 721 (1963). Schmidt, K. F., Hildebrandt, F. and Krehl, L., Klm. Wachachr., A, 1678 (1925); Knoll, A. G., Chemische Fabriken. Ger., 537, 739 (1928). Chem. Abstr. Harvill, E. K., Roberts, C. W. and Herbst, R. M. g, r . Chem., 15, 58 (1950). ' L? . Chapman, N. B., McCombie, H. and Saunders, B. C., Q, Chem. Soc., 929 (1945). D'Itri, F. M., Masters Thesis, Mich. State Univ., East Lansing, Michigan, 1966. Golton, W. C., Masters Thesis, State Univ. of Iowa, Iowa City, Iowa, 1959. Holm, R. D., Masters Thesis, State Univ. of Iowa, Iowa City, Iowa, 1958. Wehman, T. C., Doctoral Thesis, Mich. 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