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Attempted methylation with diazomethane. Diazomethane was made by adding 2.0 g. of N-nitro-N-methyl-"-nitroguanidine in portions to 20 g. of 50% potassium hydroxide solution covered by 50 ml. of ether a t 0'. The ether layer was removed, washed with water, and placed in a 200-ml. Erlenmeyer A ask. Two hundred and fifty mg. of I-hydroxy2,4-di-t-butylphenazine was added and the solution was allowed to stand overnight. The ether was evaporated and the orange crystalline residue melted without further purification a t 169.5 to 170.5'. This material was recrystallized from a chloroform-methanol mixture. It melted sharply at 170.5'. A mixture melting point with l-hydroxy-2,4di-t-butylphenazine showed no depression. Methylation with methyl sulfate. Using the method described in Organic Syntheses3 for the preparation of pyocyanine, 200 mg. of the phenazine was warmed a t about 110" with 2.0 g. of methyl sulfate for 1 hr. The reaction mixture was pouied into a mixture of 10 cc. HzO, 10 cc. methanol, and 1.5 ml. of 6N potassium hydroxide. The resulting dark solid was removed, and dissolved in alkaline methanol to a blue solution, as contrasted to the purple of the starting phenazine. The blue solution was extracted mith ether; the color transferred to the ether phase. Gradually the color changed, through green to yellow. The yellow solution gave a faint olive green solution in alkali, and a more intense red in acid. Evaporation of the yellow ethereal solution gave a gum which solidified on standing. This was taken up in boiling methanol, filtered through charcoal, and allowed to crystallize a t -10". An orange, crystalline product was obtained, m.p. 268-169", which was starting material (20 mg.). The bulk of the product, obtained on dilution with water, was a yellow gum which could not be crystallized.
WESTERNUTILIZATION RESEARCH AND DEVELOPMENT DIVISION AGRICULTURAL RESEARCH SERVICE U. S. DEPARTMENT OF AGRICULTURE ALBANY, CALIF.
An Observation on Chlorination of Normal Hexane with Iodine Monochloride' ALEXANDER I. P O P O V
AND MTILLI.4M A
Normal hexane is usually considered to be a rather stable solvent, not easily susceptible to halogenations. According to the literature, it acts as a typical "purple" solvent for halogen solutions and the solutions of iodine monochloride in this solvent should have a red-brown color with an absorption m a x i m u m in the vicinity of 460 mp. I n the course of study of iodine monochloride complexes with various Lewis bases, it was decided to use normal hexane as the reaction medium. It was immediately discovered that contrary to the expectations, dilute solutions of iodine monochloride in this solvent had a distinct purple color. Absorption spectra of these solutions were obtained and the resulting absorption curves showed a maximum ( 1 ) Paper XI1 in the series "Studies on the Chemistry of Halogens and of Polyhalides." Previous paper, J . Am. Chem. Soc., 77,4622 (1957). (2) N. N. Greenwood, Rev. Pure and Appl. Chem., 1 , 89 (1951).
22
of 520 mp which is characteristic of iodine solutions in nonpolar solvents. Although it has become evident that n-hexane was not a suitable solvent for the study of iodine monochloride complexes, it was of interest to the authors to establish if the reaction actually occurred between iodine monochloride arid n-hexane rather than with a reactive impurity in the solvent. When iodine monochloride is added to the purified n-hexane, it dissolved rapidly in the solvent, but the process of dissolution is accompanied by the evolution of a small quantity of a gaseous product, presumably, hydrogen chloride. Although in some instances, the resulting solutions did have, originally, a reddish-brown color, the latter very rapidly turned to purple and showed the characteristic iodine absorption band. I n order to see if we have a photochemical halogenation reaction, iodine monochloride solutions were prepared in a photographic dark room and the absorption spectra were obtained without any previous exposure of the solutions to illumination. I n all cases the absorption maxima were shifted toward the iodine peak. The disappearance of halogen in the n-hexane solutions was followed by preparing standard solutions of iodine monochIoride in this solvent and titrating the total halogen iodometrically in aliquots of the solutions withdrawn after definite intervals of time. The results of this study are shown in Table I. Since it has been shown that the reaction between iodine monochloride and n-hexane can occur in the dark, this series of experiments was done under ordinary illurnination. TABLE I TITRATION OF IODINE MONOCHLORIDE IN NORMAL HEXANE
DESXIN
Received May 20,1957
VOL.
Experiment I Time, Normality hr. x 103 0 1 3 5.5 23.5 100
3.26 2.64 2.12 1.75 1.66 1.66 1.65'
Experiment 2 Experiment 3 ?%me, Normality Time, Normality hr. X lo3 hr. X lo3 0 I 2 18 45 69 90
5.20 4.66 4.09 2.68 2.45 2.37 2.35
0.3 1.2 19.2 44.7 98.5 260.
5.57 3.68 2.81 2.75 2.60 2.56 2.aa
2 . 30a a Normality as calculated from absorbance data assuming only 1 2 remaining.
It is interesting to note that the limiting concentration of the titratable halogen is either equal to half, or less than half of the original halogen present. After the solutions reached an apparent equilibrium, their adsorption spectra were obtained. In all cases the remaining halogen was the iodine and the calculation of the final concentration from the absorption data agrees well with the titration value. Since there still remained a possibility that notwithstanding an apparently careful purification of
DECEMBER
1957
the solvent, some trace impurities may have remained and were responsible for the halogenation reaction, a large quantity (approx. 50 g./liter) of iodine monochloride was added to previously purified n-hexane. The mixture was allowed to stand for several days, the remaining iodine monochloride was destroyed with stannous chloride, and the solvent was dried and fractionated. This solvent was used to prepare new solutions of iodine monochloride; however, no change in the behavior of the system was observed. These results indicate that the presence of an impurity is not responsible for the halogenation reaction. Another portion of purified n-hexane was treated with a large amount of iodine monochloride for several days after which excess halogen was destroyed as described above. The resulting solution was concentrated by fractional distillation and an infrared spectrum was obtained of the residue. There was a weak but a definite indication of the carbon-chlorine stretching vibration at 700 cm.-l Unfortunately, this band was largely masked by the carbon-hydrogen rocking vibration at 720 cm.-l A solution of chlorine was prepared in n-hexane and the change in the halogen content was followed iodometrically. Although a loss of chlorine was observed, the rate of this loss was much slower than for iodine monochloride. About 50% of chlorine still remained in solution after a 12-hr. period. On the other hand, iodine solutions in this solvent appear to be perfectly stable. Solutions of iodine monochloride were likewise prepared in purified n-pentane, n-heptane, isooctane, and cyclohexane. While the first three solvents behaved similarly to the n-hexane, it was found that iodine monochloride solutions in cyclohexane appeared to be considerably more stable, although in general one would expect a cyclic hydrocarbon to be more reactive than an aliphatic one. These results agree with the observations of Buckles and Mills3 who found it possible to obtain the absorption curve of iodine monochloride in cyclohexane, provided that fresh solutions are used. These authors report an absorption maximum of 466 mp with a molar absorbancy index of 165. The experimental evidence obtained in this investigation indicates that iodine monochloride is a very active chlorinating agent for saturated aliphatic hydrocarbons. It seems to react much faster than elemental chlorine and it would consequently seem rather unlikely that the chlorination is preceded by the dissociation of iodine monochloride into molecular iodine and chloriia. It is also interesting to note that iodine is often used as a catalyst in the chlorinations of saturated hydrocarbons. It appears to be quite likely that this catalytic ac(3) R. E. Buckles and J. F. Mille, J . Am. Chem. Soc., 76, 4845 (1954).
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tion is due to the intermediate formation of iodine monochloride. A more detailed investigation of the iodine monochloride reactions with saturated aliphatic hydrocarbons would be of interest. However, such an investigation is not contemplated by the present authors as it is beyond their immediate interests. EXPERIMENTAL
Zodine monochloride. Iodine monochloride was prepared by the method of Cornog and Karges.4 It was first purified by sublimation in a desiccator over phosphorus pentoxide and the obtained product was transferred, in a dry box, to a glass apparatus consisting of a series of bulbs attached to a manifold. The apparatus was evacuated and sealed and the iodine monochloride was fractionally crystallized a t least five times. The purified product was transferred to a series of glass bulbs which were sealed until use. The m.p. was 27.2” (litU4 27.2’). n-Hexane. n-Hexane was a Mateson, Coleman, and Bell product and was originally purificd by vigorously shaking it with 10% by volume portions of fuming sulfuric acid until the acid layer remained colorless. The solvent was then repeatedly washed with dilute sulfuric acid and with water. This treatment was followed by repeated washing with alkaline permanganate, with water, and drying with barium oxide. The solvent was finally refluxed over phosphorus pentoxide and fractionally distilled through a I-meter helices-packed column. The purified product had an absorbance of less than 0.01 units, when measured in 5.00cm. silica cells a t 220 mp, which was superior to the best commercial grade of “research” grade n-hexane. The absorption measurements were done on a Cary recording spectrophotometer, Model I1 in the ultraviolet and visible regions of the spectra, and on a Perkin-Elmer infrared spectrometer, Model 13.
Acknowledgment. The authors are indebted to the Research Corp. for the support of this work and to Drs. R. E. Buckles and W. B. Person of this laboratory for many helpful discussions. DEPARTMENT OF CHEMISTRY OF IOWA STATEUNIVERSITY IOWA CITY,IOWA (4) J. Cornog and R. A. Karges, Inorganic Syntheses, vol. I, McGraw Hill, New York, N. Y., 1939, p. 165.
Detection of Flavanones by Reduction with Sodium Borohydride1 ROBERTM. HOROWITZ Received May SO, 1967
Flavanones may be conveniently redu.ced under mild conditions by sodium borohydride in aqueous or alcoholic solution. Thus, Pewz has reported the (1) (a) Presented in part before the 131st Meeting of the American Chemical Society, Miami, Fla., April 1957; abstracts p. 58 L. (b) A method which is closely similar has since been reported by E. Eigen, M. Blitz, and E. Gunsberg. Arch. B i o c h a . and Biophysics, 68, 501 (1957)(2) J. C. Pew, J . Am. Ghem. SOC., 77,2831 (1955).
