Aug. 20, 1959
EVIDENCE FOR
THE
EXISTENCE OF
PEROXYURANIC
ACID
4167
compound. Pcrflu,Jrocyclopciitaiie was used as a reference melting point, infrared spectrum and mass spectrum were material with a field strength of 15,020 gauss. T h e S308F2 identical with those found for the product obtained from the showed a single peak at a negative shift of 31.7 p.p.m. from reaction of sulfur dioxide with pcroxydisulfuryl difluoride. A the reference sample. Comparison of areas under the curves second product of this reaction was identified as pyrosulfuryl corrected for different densities gave approximately a 5: 1 fluoride, S205F2. ratio for the number of fluorine atoms per molecule. These Structure.-From the fact t h a t one mole of S308F2reacts d a t a show t h a t the compound S308F2 has two fluorine atoms with a solution of potassium hydroxide to give approximately per molecule, both having the same chemical environment. two moles of S O J - a n d one of Soap-,Lehmann and Kolditz proposed the structure Mass Spectra.-Mass spectra were obtained with a Consolidated Electrodynamics Corp. Mass Spectrometer, Model 000 21-103C with an ionizing current of 10.5 pamp. and an ionizFSOSOSF ing potential a t 70 volts. T h e butane sensitivity was 27.7 000 divisions per micron. Spectra were obtained for S3OsF2 as well as for S20~F2and S206F2 for comparison. The S I O ~ F ~This structure is consislent with the iiuclcar inagiletic rcsospectrum showed major peaks corresponding to SOzF +, tiancc spcctruxn which indicates only one type of fluorine SOF+, SOz+, S03+ and numerous minor peaks. The parent atom, but it is not consistent with a rnitss spectrum eontainpeak a t inass number 262 was missing in all spectra. Pering a S:05F2+ peak. The writers are of the opinion t h a t the sistent pcaks at mass numbers 182 and 184 corresponding above structure is correct and t h a t the S ~ O A F Z peak + in the to S?05F2 + could not be explained, companion with the specmass spectrum is probably due t o S?OSFL as indicated carlier. trum of S20SF,showed marked similarities even though all The name trisulfuryl fluoride was assigned to S308F2by other tests showed the absence of S206F2 in the S308F2 the discoverers. If the above structure is correct the name samplc. It is probable, therefore, t h a t the S308F2 molecule sulfuryl fluorosulfonate is equally good. is extensively disintegratcd in the mass spectrometer or in the Acknowledgment.-This work was supported in line for admitting the gas to the spectrometer, with the production of Sz05F2. part by the Office of Kava1 Research. Preparation from SO,and BF3.-Following the method of WASHINGTON Lehniann and Kolditz3 a cornDound was isolated which had a SEATTLE, molecular weight of 267 as determined by gas density. T h e AMHERST. MASSACHUSET rS
[CONTRIBUTION FROM
THE
NATIONAL LEADCOMPAXY OF OHIO]
Evidence for the Existence of Peroxyuranic Acid1 BY R. E. DEMARCO, D . E. RICIIARDS, T. J . COLLOPY AND R. C. ABBOTT RECEIVED MARCH 16, 1959 Uranium tric.xidc \vas trcatcd with aqueous hydrogen peroxide to form a hydrated uraniuni percisidc ctiinpound identical to t h a t precipitated from aqueous uranyl nitrate sdution by the addition of hydrogcn peroxide. The stoicliioinctric conipound, U 0 4 . 2 H 2 0 ,was shown t c react with alcoholic potassium hydroxide to form hydrated monobasic ( K H U 0 5 , x H 2 0arid ) dibasic ( K Z U O 5 d - I 2 0 potassium ) salts. A hydrated peroxyuranic acid structure (H2LT05.H,0) was concluded to best represent a compound of the sttrichiometry U04.ZH20. This structure w11s supported by the infrared absorption spectrum and reactions with a strong base cf the uranium peroxide compound.
Introduction
In light of the recent work by Gentile, Talley and Collopy,8 further substantiating the acidic Since Fairley? first prepared hydrated uraniuni character of hydrated uranium trioxide by its reperoxide by the precipitation of uranium from action with the weak base urea to form urea uranate, weakly acidic aqueous solutions of uranyl nitrate additional study of the reactions of the peroxyor uranyl acetate with hydrogen peroxide, a con- uranium hydrate compounds was warranted. troversy has existed regarding its structure. DePreliminary tests employing uranium peroxide spite the controversy regarding its structure, the hydrates and aqueous or alcoholic solutions of stoichiometric compound U04.2Hz0 readily can be urea failed to produce sufficient quantities of a prepared with a high degree of purity. compound corresponding to urea peroxyuranate. Previous experimental efforts to define the struc- A solid material corresponding to urea peroxyurature of this uranium peroxide compound were based iiate was precipitated, however, from a concenon its reactions in aqueous solution, its stepwise trated solution of diaquotetraureadioxouraniumthermal decomposition to Us08 and/or infrared (VI) nitrate by the addition of hydrogen peroxide. spectral studies. As a result of these experiments This precipitate, after continued contact with water, the following structures have been postulated for hydrolyzed to a peroxyuranic acid hydrate. the compounds: U04.2H20,3UOz(02).2H20,4 LOa. In view of the above results and the generally H?OZ.H~O HzUO5.H2O6 ,~ and H4U06.’ diminished acidity of peroxyacids as compared with fundamental oxygen acids of the same system, a (1) T h i s paper is based on work performed for t h e Atomic Energy series of tests was conducted with the strong base Commission by t h e National Lead Company of Ohio a t Cincinnati, potassium hydroxide in alcoholic media to demonOhio. (2) T . Fairley, Chem. Y e w s , 33, 237 (1876). strate the acidic character of uranium peroxide (3) A. Rosenheim and H . Daehr, Z. anorg. allgem. Chem., 181, 177 hydrates. These experiments were supplemented (1929). with infrared spectral data to determine the pres(4) G. Tridot, Compt. r e n d . , 232, 1215 (1961). ence of the uranyl group, hydrate water and hy(5) G. F. Hilttig a n d E. v. Schroeder, Z . anorg. allgem. Chem., 121, 243 (1922). droxyl groups. X-Ray diffraction patterns were (6) A . Sieverts and E . L. Miiller, ibid,,173, 207 (1928). ( 7 ) J J. Katz and E. Kabinowitch, “ T h e Chemistry of Uranium.” LIcGrair Hill Book Cm, I n c , New Y o r k , S . Y . , 1051, pp. 245-246.
18) P S Gentile, I, H Talle\ a n d T J Collopy, J 10, I14 (1950)
(hem
Znnrg A i d e a r
obtained to verify the existence of a single homogeneous product phase. Experimental Materials.-I. The hl-drated uranium peroxide cornpound was precipitated from an aqueous solution of uranyl nitrate hexahydrate (Reagent Grade purchased from Baker Cheniical Company) using the procedure described by W a t t , Achorn and hlarley.' The resulting precipitate was filtered, washed with distilled water arid then alcohol, prior to drying a t 100' for 60 minutes. Calcd. for H?UOj.H&: U, 7U.41; peroxy ox).geii, 9.46. Found: U , 70.29 k 0.07; peroxy oxygen, 9.41 =k 0.11. 2. Tlic brick-red X-ray amorphous form of urmiuin trioxide, prepared by the thermal decomposition of U 0 3 , 2Hz0 (Calcd. U, 73.91. Found: U, 73.90 i 0.07) a t 400" mas slurried with an excess of aqueous 15';; H.02 solutioli at rooni temperature for 20 hr. The product of this slurry was filtered, washed with distilled water and tlicri a1COhlJ~,prior to drying itt 100"for 60 minutcs. Calcd. for H?UOj.HzO: U, 70.41; peroxy oxygen, 9.4G. Found C , 70.50 =k 0.07, peroxy oxygen, 9.44 + (1.11. Analytical Methods.-Uranium, potassium and peroxy oxygen were determined in accordance with the procedures given by "Scott's Staiidard Methods uf Chemical Analysis. " l o Lraniurn was determined (1) gravimetrically by ignition to U30, or (2) volumetrically by titration with standard 0.1 N potassium permanganate. The volumetric rnetliod of analysis was used for all potassium containing samples. Potassium was determined gravimetrically by precipitation as potassium perchlorate from an ethj-l alcohol-butyl acetate solution. Peroxy ox?-gen was determined by dissolution o f the sample in dilute sulfuric acid and titratioii with standard 0.1 N potassium permanganate. lVater was determined by titration of methyl alc(~Iiijlslurries o f the sample with Karl Fischer reagent." Infrared spectra were obtained with a Beckmaii Motlel Ill-4 spectro~neterusing putassium bromide pellets containing 1 w/o ~arnple.1~>13 S - R a y powder diffraction patterns were obtained a t room tcinperature with a Norelco X-raJ- diffraction unit using 114.7 ~ n m cameras . and uuresolved Cu Kal-a2 radiation. Procedure .--Fivegram samples of the hydrated uranium peroxide compound (preparations I and 11) were treated with 200 nil. of itri alcohol solution of carbonate-free potassium hydroxide ranging in strength from 0.02 to 4.0 molar. The reaction was allowed to proceed, a t room temperature with vigorctus stirring, for 24 lir. before the product was filtered, alcohol-washed free of excess base and air-dried t o a constant weight.
of hydrate water16 in the starting peroxide compound. The broad infrared absorption band with maxima a t 2.94 and 3.20 I.( is indicative of the presence of hydrogen-bonded hydroxyl groups. The remaining absorption band a t 11.05 may be attributed to the asymmetric vibration of a uranyl group, one oxygen of which is hydrogen bonded to a water molecule. Hydrogen bonding to the oxygen of a M=O group usually results in a frequency shift similar in direction and magnitude to the 30 cm.-' shiftl7ml8observed in the spectrum of the uranium peroxide compound. h summary of the analytical results (and the mole ratio values calculated from these results) for the reaction of either uranium peroxide hydrate starting compound with alcoholic KOH solutions is given in Table I. TABLE
ACTION KOH,
M
0.02 .10 .50 2.0 4.0
-Chemical
u
67.60 61.51 58.90 57.15 52.56
K
I
HpUOs'HpO WTTH KOH analyses, (,'hP X I o l e ratici--.
2.50 11.07 13.66 15.34 18.07
OF
0-0
H20
U
9.19 7.90 7.66 7.50 6.96
10.7 10.3 11.1 12.7 11.8
1.00 1 .oo 1.00 1.00 1.00
K
H20
0-0
0.23 1.01 1.00 0.96 L41 .97 1.63 .98 2.10 ,$I9
2.1 2.2 2.5 2.9 3.0
X-Kay diffraction patterns of the starting peroxide compound and of the potassium salts formed were used to supplement chemical analysis by ascertaining the existence of a homogeneous product. The product of reaction with 0.02 M KOH was, by this means, identified as a mixture of HPU05.Hz0 and hydrated KHUOb. Similarly, the products of reaction with 0.5 and 2.0 31 KOH were identified as a mixture of hydrated KHUOs and KzU05. The large excess of base required for the formation of the hydrated monobasic (KHUO5.2HaO) and dibasic (KzU05.3Hz0) salts in alcoholic potassium hydroxide solutions can be explained by the ready hydrolysis of these salts to HzU05.HzO in the presence of water. This hydrolytic tendency explains the failure of previous investigators to Results and Discussion establish the acidic nature of the uranium peroxide Infrared spectra and X-ray diffraction patterns compound via a simple acid-base reaction in of the two uranium peroxide starting compounds aqueous media.IY,?O Hydrated potassium peroxywere obtained. The infrared spectra and diffrac- uranates have been prepared, however, by other tion patterns of these materials were not affected methods not designed to establish the acidic by the mode of preparation. Strong infrared ab- character of H2U0h.H20. Formation of the hydrated monobasic salt, sorption bands were found a t 2.94, 3.20, (3.18, 10.65 and 11.05 p ; a shoulder occurred at 11.55 I.(. KHU06.2Hz0, and the hydrated dibasic salt, The absorption band a t 10.65 h and the shoulder KPUO5.3H20,was verified by chemical analysis, a t 11.55 p previously have been associated with the X-ray diffraction techniques and a study of the asymmetric and synimetric vibrations of the uranyl thermal decomposition reactions of these salts. g r o ~ p . The ~ ~ band , ~ ~ a t G.15 p is due to the presence The monobasic salt was shown to decompose to potassium pyrouranate on heating in accordance (!IG ). W.W a t t , S . L , Achorn and J. L . Llarley, THIS J I J U R N A I . , 72, with the equation 3311 (1930). (10) N . H. Furman, E d . , "Scott's Standard Methods of Chemical Analysis," D. Van Nostrand Co., New York, N. Y.,1939, p p 871, 1022, 1027, 2180. (11) J . hlitchell, Jr., and D. M . Smith, "Aquametry," Interscience Publishers, Inc., Kew E'ork, N Y . , 1918, p p . 02-102. ( 1 2 ) 11,R l . Stimson and X I . J. O'Donnell, THISJ I J C J R ~ - A 74, T . ,1803 (1Y.52) CW 270S(195'2). C~., (13) U. Schiedt a n d H. Reinwein, Z . S Q ~ Z W ~7b, (1.1) G. K. T . Conn a n d C . K . W u , 2'roiis. P a r a d a y .Yoc., 3 4 , 1483 (1038). 1 1 . j ) (', 'l'riO H2S05,HzO Hydrated peroxyoxygen acid H2COj.H~O H&O? P5ro (oxygen) acid H2U20i
The preceding data and discussion were primarily intended t o establish the structure of hydrated uranium peroxide through reaction with strong bases. I n light of the previously established amphoteric character of other hexavalent uranium oxide hydrates, peroxyuranic acid would be expected to react, as a base, with strong acids. Acknowledgment.-\Ye wish t o express our appreciation to F. H. Ford for obtaining and interpreting the X-ray diffraction patterns, to B. Gessiness and his staff for the analytical results, and to P. S. Gentile, L. H. Talley and D. A. Stock for their help in obtaining the above data. CIXCISSATI,OHIO
Low TEMPERATURE LABORATORY, DEPARTMENTS OF CHEMISTRY A N D CHEMICAL ENGINEERING UNIVERSITY O F CALIFORNIA, BERKELEY]
The Low Temperature Heat Capacity and Entropy of Thallous Chloride1 BY I. R. BARTKYAND W. F. GIAUQUE RECEIVEDFEBRUARY 19, 1959 The heat capacity of thallous chloride has been measured from 15 to 310°K. The entropy at 298 15'K was found to be 26.59 cal. deg.-' mole-'. T h e heat capacity, entropy, (P- Hoo)/Tand ( H O- Ho0)/Tfunctions for TlCl are tabulated from 15 t o 300°K. The entropy change calculated from the third law of thermodynamics for the cellreaction T1 AgCl = TlCl Ag was found to be in good agreement with the cell temperature coefficients of Gerke provided the entropy change for the step Tl(s) = Tl(amalg) is calculated from A S = (-lHLalorlmotrlo FE)/T rather than from dE/dT. It is evidently difficult to obtain reliable cell temperature coefficients involving a metal electrode even when it is as soft as thallium, Third XgCl = TlCl -Ig, AH0 = law results have been used to calculate these heats of reaction a t 298.15"K.: Tljcryst.) -18,430 cal. mole-'; T1 l/1C12 = TlC1, AHu = -48,800 cal. mole-'.
+
+
+
+
+
Gerkea studied a number of cells in order to test the validity of the third law of thermodynamics by comparing the entropy changes as calculated from the temperature coefficients of the cell voltages with those calculated from the available low temperature heat capacity data by means of the third law. He found a discrepancy of 2.0 cal. deg.-l mole-' in the cell T1
+ .igC1 = TlCl + -Ig
(1)
The heat capacity data available at that early date often were inaccurate since the techniques of low temperature calorimetry were in the first stages of their development. Thus Gerke considered that the discrepancy had its origin in the heat capacity data and particularly in t h a t on thallous chloride. The present work provides accurate heat capacity data for thallous chloride. Calorimetry and Material.-The low temperature heat capacity was measured in 311 apparatus similar to that described by Giauque and Egan . 3 T h e particular caloritn(1) This work was supported in part by t h e National Science Foundation. T h e U. S. Government may reproduce this article. (2) R . H . Gerke. THIS JOURNAL, 4 4 , 108.1 (1922). (:i) W. IT. Giauclue and C. J . Bgan, J. C h e i n . Phys.. 5 , 4.5 (I'JH7).
+
eter was described by Kemp and Giauque4 and its most recent modification by Papadopoulos and Giauque.6 Briefly, the present calorimeter was of copper and the temperature was measured by a gold resistance thermometer-heater, for high precision, and Laboratory Standard Copper-Constantan Thermocouple hTo. 105 for simultaneous reference before and after every heat input. The thermocouple was compared with the triple (13.94'K.) and boiling points (20.36' K . ) of hydrogen and the triple (63.15OK.) and boiling points (77.34'K.) of nitrogen, during tlie course of the present work. O D was taken as 273.15"K. and 1 defined calorie was taken equal t o 4.1840 absolute joules. The thallous chloride was supplied 99.9+'r, pure by the Chemical Commerce Co. It was received in tlie form of a rather voluminous precipitated powder. This material was heated to the melting point (about 430°), in a Pyrex beaker so t h a t a large reduction in volume occurred. This eliminated the possibility t h a t particles with microscopic properties could remain in the sample used for the heat capacity measurements as could be the case with a fine powder. The well-crystallized solid was broken easily with a mortar and pestle and the crystals were passed through an 8 mesh i n - l screen so that they would pack well within the calorimeter. The use of this material, rather than the powder, a t least doubled the amount of material which could be placed in the calorimeter and a1m improved the heat con(4) J. D. Kemp and W . F.C.iauque. THISJ O U R N A L , 59, 70 (lW37). (3) hl. N. Papadripoulos and W. F. Giauque. ibid., 7 7 , 27-10 (1955).
