August, 1962
MOLECULAR STRUCTURES OF INPERHALOGEN
COhlPOUKDS
1397
MOLECULAR COMPLEXES OF SOME IKTERHALQGES COMPOUNDS1 BY MAXT. ROGERS AND W. K. MEYER~ Kedxie Chemical Laboratory, Michigan State University, East Lansing, Michigun Received January 19, l Q S 8
Molecular complexes of iodine trichloride, iodine pentafluoride, iodine monochloride, and iodine monobromide with various electron donors, such as dioxane and derivatives of pyridine, have been prepared and some physical properties of the more stable ones have been studied. Dissociation constants of some of these complexes have been determined spectrophotometrically and dipole moments have been estimated from the dielectric polarizations of their solutions in non-polar solvents.
Introduction
Experimental
Materials.-Benzene (d2$ 0.87336), carbon tetrachloride Stable molecular complexes between iodine, (d*$ 1.5844), and dioxane (m.p. 11.66') were purified by iodine monochloride, and iodine monobromide and partial freezing and distillation over a drying agent. The various electron donors such as pyridine and diox- pyridine and pyrazine derivatives were commercial materials A sample of pyrazine was the gift of the Wyanane long have been known and the st,ability and redistilled. dotte Chemical Corp., Wyandotte, Michigan. 2-Fluorovarious other properties of a number of these have pyridine and 3-fluoropyridine were prepared from 2-aminobeen intensively investigated r e ~ e n t l y . ~Appar- pyridine and 3-aminopyridine by diazotization in 40% ently the only stable complex of iodine penta- fluoroboric acid.9 Iodine monochloride (m.p. 2 i 0 ) , iodine trichloride, and fluoride reported is a 1:l complex with d i ~ x a n e . ~iodine monobromide (m.p. 41.5') were prepared from the Some complexes of iodine trichloride have been elements and (except for iodine trichloride) were purified but no physical measurements appear by recrystallization from the melt. Iodine entafluoride to have been carried out on them. We therefore from the General Chemical Division, Allied Zhemical and Dye Co., was fractionally distilled in t-i Monel-fluorothene have attempted to prepare EL variety of complexes stilPO beforeuse (m.p. 9.6'). of the interhalogen compounds using as electron Preparation of Complexes.-Normally a solution of the donors pyridine and pyrazine and some of their interhalogen compound in carbon tetrachloride was added derivatives, dioxane, trifluoroacetic anhydride, to a solution of the electron donor, also in carbon tetrachloand the crystalline complex was removed by filtration, amines, ethers, etc., and to study some of their ride, washed with carbon tetrachloride, and dried. physical proper ties. Dioxane-iodine monobromide and dioxane-iodine pentaIt is particularly useful to know the extent to fluoride were prepared by mixing the pure liquids in stoichioamounts, filtering, washing with carbon tetrawhich these complexes dissociate in solution to metric chloride, and drying. Some of the complexes were liquids give the original materials and we therefore have at room temperature. measured tqectrophotometrically the dissociation Known weights of pure iodine pentafluoride were dispensed constants of some of them. Since the dielectric from B calibrated 2-ml. Pyrex buret fitted with a stopcock a Teflon plug. All transfers were carried out in an constants of the solutions lead to estimated dipole having atmosphere of dry nitrogen. moments for the molecular complexes when the Analyses.-Microanalyses of some complexes were carried dissociation constants are known, we have measured out (for C, H, N ) by Spang Pvlicroa8nalyticalLaboratories, dielectric constants of a number of the crystalline Ann Arbor, Michigan, and are reported in Table I. Halogen analyses were made for several complexes by complexes dissolved in non-polar solvents. From determining an iodometric equivalent by the following the electric moments of the complexes the extent technique developed for these compounds. of electron transfer from donor to acceptor was A weighed sample (2-3 meq.) of complex was dissolved estimated. The electric moments of the dioxane- in 10 ml. of pyridine, 6 gbof potassium iodide was added, mixture was cooled to 0 , and 12 ml. of 12 N hydrochloric iodine and pyridine-iodine complexes in cyclo- the acid was added. Acetone (10 ml.) was added to dissolve hexane solution have been measured previously and the liberated iodine and standard thiosulfate was added the values 3.0 and 4.5 D., respectively,7 reported; until the mixture was yellow. Then 20 ml. of water and also, the moment of the pyridine-iodine complex 2 ml. of starch solution were added and the titration was cont,inued to the end-point. A blank also was run. in benzene solution*has been reported as 4.17 D. The calculated iodometric equivalent is taken to be the We also have obtained parts of the phase dia- molecular weight of the complex divided by the number of grams for the two-phase solid-liquid systems diox- halogen atoms considered to be present. The observed values for the iodine trichloride complexes are somewhat ane-iodine peiitafluoride and pyridine-iodine pen- high, corresponding to 5-10% loss of chlorine from the tafluoride. complex. Dielectric Constants.-Dielectric constants of a series of (1) This work was supported by a grant from the Atomic Energy six or seven solutions of the molecular complex in benzene Commission. or dioxane, usually varying from 0.0005-0.02 mole fraction of complex, were measured using a heterodyne-beat ap(2) Abstracted from a thesis submitted by W. K. Meyer in partial paratus; densities were measured with a pycnometer. fulfillment of the requirements for the Ph.D. degree, Michigan State The methods of measurement of dielectric constant and University, 1958. densit as well as the methods of calculation of results have (3) A . I. Popov, C. Castellani-Bisi, and Willis B, Person, J . Phys. been rfescribed.11 The plots of the dielectric constants and Chem. 6 4 , 691 (1960), and preceding papers of that series. densities of the solutions vs. mole fraction of solute were (4) A. F. Scott and J . F, Bunnett, J . Am. Chem. SOC.,64, 2727 linear within experimental error in every case. Molar (1942). refractions were calculated from tables of a t o r i c refractions. ( 5 ) E. V. Zappi and M. Fernandez, Analrs Asoc. Quim. Arg., 27, Phase Diagrams.-Freezing points of liquid mixtures 102 (1939). were measured in a small Pyrex dewar vessel in which the (6) R . Cernstesou and M. Poni, Analele Acad. Rep. Populave Romane, Ser. Mat. Fie. Chirn.. 8 , 140 (1950). (7) G . Xorttim and H. Walz, 2. Elektrochem., 57, 73 (1953). (8) Y , I
