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Vol. 73
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0.4-3.0 M , radioactive assays after 6 days showing low solu- performed by dissolving 1.5 g. of Analar hydrated bilities of about 3 mg. Am/liter. However, the rate of pre- sodium sulfite [NazS03.7H20) in 3 ml. of enriched cipitation after initial oxidation was dependent upon carbonate concentration. For example, 4.5 hours after oxida- water a t room temperature ( ss 18’). This solution tion in 0.4 M potassium carbonate solution, the americium was exposed to oxygen, confined in a gas buret, concentration remaining in solution was -60 rng./liter, until no further absorption of gas took place: while in 3.0 M potassium carbonate the unprecipitated in all cases the absorption was close to the theoamericium was -23 mg./liter. It was found possible t o dissolve quickly the insoluble retical requirements. That oxidation was complete americium compound in 0.1 ilf sulfuric acid after washing was checked by titration with KMn04 and also by with distilled water. The solution appeared colorless and gravimetric estimation of the sulfate formed. while being mFaswed in the spectrophotometer, the absorp- The 01*content of the water was measured before tion at 5030 A. was checked a t intervals as an index of the and after the oxidation by determination of the Am(II1) present. During the course of the measurement (2.5 hours), the Am(II1) increased from an initial concen- water density, after vigorous purification, by the tration of 2% t o a final value of 5y0. By correcting the semi-micro silica float method.3 The results are light absorption due t o Am(III), the curve shown in Fig. IC summarized in Table I : i t is clear from experiments was obtained which is taken to be characteristic of Arn(V) 4 and 5 , where the solution was left for some 72 in 0.1 M sulfuric acid solution. The absorption bands are time longer than a t 5150 and 7200 A. with molar extinction coefficient 48 and hours a t room temperature-a required for the completion of oxidation in experi66, respectively. Final proof of the oxidation state of the new form was ob- ments 1 , 2 , 3 and 6-that the exchange of the sulfite tained by titration with a reducing agent. In a typical ion is negligible under these conditions, the obexperiment, americium was oxidized and precipitated as described, washed six times with distilled water, and dis- served decrease in water density in p.p.m., Ayd, solved in 0.1 M sulfuric acid. I t rvas precipitated again being essentially that to be expected from dilution by addition of potassium carbonate, washed four times again with the water of crystallization of the sodium with water and dissolved in 0.1 -11 sulfuric acid. From this sulfite. Experiments 1, 2, 3 and 6, show conclusolution an aliquot of 0.100 ml. was removed and the con0 2 --+ centration measured10as 7.24 * 0.03 millimolar. Measure- sively that the over-all reaction 2NazS03 ment of the absorption qpectrum showed 4cI of the ameri- 2NazS04 proceeds without exchange of oxygen with cium to be Am(II1). the solvent water. Exactly 0.030 ml. of 0.0502 .V ferrous ion solution was TABLE I added, atid as well as could be told by spectral analysis, A.yd reduction to Am(II1) ~ w scomplete. The solution was Vol. of O z adsorbed, calcd. for transferred quantitatively t o a micro titration dish and the ml. N . T. P. Ayd 7H10 from Experimental excess ferrous ion titrated with standard ceric solution using Expt. Obs. Calcd. obs. sulfite conditions ferrous phenanthroline as an indicator. It rvas found that Oxygen 1 66.0 66.6 -79 -78 1.367 microequivalents of ferrous ion had been oxidized by from -81 -78 2 66.2 66.6 the americium which upon calculation s h o w d the oxidized state to have 1.96 equivalents/mole. 3 66.8 66.6 -80 -78 cylinder Since the forms of Pu(V) and ;\;p(V) ions in acid solution 4 .. .. -80 -78 In atmosphere have been shown11 to be PuOzf and NpOl-, respectively, of nitrogen we may expect Am(V) t o be Am02+undersimilar conditions. 5 .. , . -79 -78 In vacuo N o attempt has been made to identify the Am(V) compound which precipitates from carbonate solution. 6 66.5 66.6 -80 -78 In oxygen, trace This work was performed under the nuspices of the 11. S. CuSO4 added AEC. Ayd calcd. for 7H20 4-0 = - 91. (10) Measured by a-particle assay using a half-life of 475 years. Ayd calcd. for 7H20 3 . 0 = - 117. (11) K.A. Kraus and G. E. Moore, National Nuclear Energy Series, Ayd calcd. for 7&0 4 . 0 = - 130.
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Plutonium Project Record, Vol. 14B, “The Transuranium Elements, Research Papers,” Paper No. 4.19 (McGraw-Hill Book Co.,Inc., New York, N. Y.,1949); L. B. Magnusson, J C. Hindman and T. J LaChapelle, ibid , Paper No, 15.4.
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Examination of the various chain mechanisms which have been proposed for this reactioninvolving, e.g., either the radical HSOs4 or the RADIATION LABORATORY AND DEP.4RTMENT O F CHEMISTRY radical ion SOP shows that on either hypothesis U N I ~ R S I TOF Y CALIFORSIA RECEIVED AUGUST31, 1950 the above result might be expected, since oxygen BERKELEY 4, CALIF. transfer involving oxygen from the water molecules, would not be expected to occur. It is however Oxygen Atom Transfer during the Oxidation of unfortunate that our experiments were of necessity Aqueous Sodium Sulfite (because of the relatively small enrichments of HzO1*available at the time) performed using soluBY E . R. S. WINTERAND H. 5’. A . BRISCOE tions very much more concentrated than those The very interesting communication of Halperin which have been used in investigations of the chain and Taube’ on the above subject prompts us to reaction. It would be very interesting to know put on record some experiments upon similar lines if these results are confirmed in more dilute soluwhich were performed in 1939.2 In this work we tions. studied briefly the oxidation by normal oxygen For experiments upon the oxidation of sodium gas and by normal hydrogen peroxide of normal sulfite with hydrogen peroxide, the exchange of sodium sulfite, dissolved in water containing 01* hydrogen peroxide with water was first examined. in excess of the normal abundance. In both cases Merck perhydrol (aqueous hydrogen peroxide the oxygen atom which was transferred to the containing 29% H202 by analysis) was used as a sulfite ion came from the oxidizing agent, and not source of hydrogen peroxide. It proved impossible from the water. (3) Briscoe, et al., J . Chcm. Soc., 1207, 1948 (1934); Spoor, Ph.D. The experiments upon oxygen oxidation were Thesis, London, 1935. (1) Halperin and Taube, THISJOURNAL, 72,3319 (1950) (2) Winter, Ph.D. Thesis, London Uuiversity, 1942.
(4) Franck and Haber, Sifsungsber. Prcsss. Akad., 50 (1931). ( 5 ) Backstrom, Z. physrk. Chcm., 26B, 122 (1934).
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Jan., 1951 to mix the perhydrol with the heavy water and subsequently separate b y physical methods the bulk of the hydrogen peroxide from the water, so that a direct check of the extent of any exchange was not feasible. Instead, 1-ml.lots of the perhydrol were mixed a t room temperature with 2-ml. portions of heavy oxygen water, and the solution left for 48 hours with addition of a trace of a catalyst active in promoting the decomposition of hydrogen peroxide. At the end of this time the liquid was raised just to the boiling point to complete the decomposition, and the residual water purified and its density determined. The complete disappearance of the peroxide was confirmed by chemical. tests. I n a further experiment the mixture of peroxide and heavy water was left for 24 hours a t room temperature before addition of the catalyst. The results are shown in Table 11, from which i t is evident that the decrease in density observed is wholly accounted for by the water produced in the decomposition of the peroxide and that added with the perhydrol, and we conclude that the oxygen evolved did not suffer exchange with the enriched water. Clearly under these conditions there is little if any exchange of the hydrogen peroxide molecule with water : neither does the heterogeneous catalytic decomposition require the active participation of the water molecule.
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sulfite. The experimental error in the water density determinations recorded here is *l r d (1yd = 1 in lo6of density). DEPARTMENT OF INORGANIC AND PHYSICAL CHEMISTRY IMPERIAL COLLEGE S. W. 7, ENGLANDRECEIVED SEPTEMBER 29,1950 LONDON,
Diphenyl-t-dehydroabietinol and its Oxidation Product BY HAROLD H.ZEISS
I n the course of continued work with the resin acids i t has become apparent to us that our earlier reported studies in this field are subject to correction. It has been found that the procedure described for the isolation of diphenyl-t-dehydroabietinol (I)' leads instead to a cleavage product, which, on treatment with hot chromic acid, is oxidized to ketonic material of different structure than that previously proposed.2 In view of the recently published paper of Brossi, Gutman and Jeger3 in which the same conclusions are reached, it is considered appropriate to report our results a t -~ this time. Authentic I is obtained from the Grignard reaction between methyl dehydroabietate- and phenylmagnesium bromide according to the original procedure, with the exception of subjecting the crude product to crystallization from glacial acetic TABLEI1 acid instead of to distillation. The ultraviolet Ayd Ayd calcd. for calcd. for spectrum of the crystalline carbinol, m.p. 139O, 1.0 from 2.0 from Catalyst Obs. Ha02 HiO, is consistent with that of the dehydroabietic-type structure as exemplified by the spectrum of deMnOz Mn208 -51 -50 -56 hydroabietic acid (Fig. l). If, however, the crude Fez08 -51 -50 -56 carbinol is vacuum distilled in a Claisen flask, a Palladized asbestos -52 -50 -56 yellow oil is collected which analyzes correctly for Platinized asbestos -52 -50 -56 I but which gives an ultraviolet spectrum similar Platinized asbestos added after to that of benzophenone. The explanation for 24 hr. -52 -50 -56 this has been secured from the observation that The perhydrol was then used to oxidize sodium crystalline I, when heated between 190-230° sulfite: 2 ml. of perhydrol was mixed with 4 ml. of and then cooled to room temperature, is transenriched water, and 1.71 g. of finely powdered formed into a yellow liquid which also gives a anhydrous sodium sulfite added slowly with stirring. benzophenone-type spectrum (Fig. 2 ) . The presThe temperature of the mixture reached 40" ence of benzophenone is confirmed by its isolation during the oxidation. A t the end of the oxidation from the mixture. the mixture was boiled, cooled and the density of The thermal cleavage of I is reminiscent of the the water determined as before. A parallel experi- type experienced by other workers with tertiary ment showed that substances capable of decoloriz- alcohols.4 In this instance the expected products ing permanganate were equivalent to less than 1 are benzophenone and nordehydroabietane (11. ml. of 0.1 N KMn04 at the end of the reaction. 12-methyl-ll2,3,4,9,10,11,12-0ctahydroretene). With the quantities used above, the decrease in H\ ,CHa density to be expected due t o the water added as perhydrol, including that which would remain after the reaction HzOz Na2SOa -+ H2O Na2S04 was Ayd = -50. The total decrease to be expected if the second oxygen atom in the hydrogen peroxide molecule CHs suffered exchange with the heavy water was Ayd = \ -56, while the total decrease to be expected if the I1 three oxygen atoms of the sodium sulfite also exchanged was Ayd = -65: the observed decrease Evaporative distillation and chromatography were in water density was Ayd = -52, demonstrating (1) H. H. Zeiss, THISJOURNAL, BB, 302 (1947). (2) H.H.Zeiss, ibid., 70,858 (1948). that, allowing for the normal error, the experimental (3) A. Brossi, H.Gutmann and 0. Jeger, Relv. Chim. Acta, 88, 1730 technique, and by comparison with the results in (1950). Table 11, direct transfer occurs of the oxygen atom (4) P. Ramart-Lucas, Ann. chim., [SI 80, 349 (1913); V. Grignard from the hydrogen peroxide molecule to the sodium and F. Chambret, Compf. rend., 182, 299 (1926). ~
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D\ G,
