1984
NOTES
Vol. 63
the sotationary arrest points was estimated to be within f2 The compositions in mole fractions are based upon the weights of the added salts.
.
800 6
700
P.. 600
500
0.4 0.6 0.8 PL I PuCls, mole fraction. Fig. 1.-Phase diagram of the binary system PuCL-KC1 as d e t e r e d by thermal analysis: (1) PuC13melting point, 769 j = 2 ; (2) eutectic point, 486 f 3" at PuC13 mole fraction 0.57; (3) peritectic point for the compound KzPuC15, 611 f 3" a t Pu$& mole fraction 0.35; (4) K3PuCls melting point, 685 f 3 . (5) eutectic point, 621 f. 3' at PuC13 mole fraction O.l$'; (6) KCl melting point, 771 f 2'.
KC1
0.2
form by transferring the melt under vacuum into a Pyrex mold and stored in a vacuum desiccator. The melting point of the potassium chloride was determined to be 771 =k 2' in agreement with the reported value.6 Hydrogen chloride (Matheson Co.) and argon (Linde Air Products) were dried with phosphorus entoxide. Apparatus.-The mezing chamber consisted of a 32 cm. long, 22 mm. diameter quartz test-tube with a Teflon stopper in which was mounted a 35 cm. long, 5 mm. quartz tube for the gas inlet and 35 cm. long, 3 mm. diameter quartz thermocouple well. The lower 20 cm. of the melting chamber was surrounded by graphite to ensure uniform temperatures. A Pt-lO% R h thermocouple which checked within f 0 . 4 " with the melting points of zinc and aluminum was employed for the temperature determinations. The thermocouple was provided with a plug-and-jack connection for the purpose of switching from a Honeywell 0-1650° recorder used to indicate the liquidns points to another thermocouple a t 0' and a K-2 potentiometer used t o measure the temperature a t the stationary arrest points. Procedure .-To begin an experiment the desired quantities (20 to 80 g. total) of salts in the melting chamber were melted and stirred with a stream of hydrogen chloride or argon gas. The thermocouple well projected to within 5 mm. of the bottom of the melting chamber and the volume of the solution was such that the thermocouple was immersed to a depth of 2.5 to 6.0 cm. After 0.5 hr. of stirring with the melt approximately 70" above the liquidus, the cooling curve was begun. Although hydrogen chloride was employed in most of the experiments, identical cooling curves were obtained with argon. The accuracy of the temperature measurements a t ( 5 ) National Bureau of Standards Circular, Number 500, 1952. p.
804.
Results The experimental results are summarized in Fig. 1. These data indicate the existence of the compound K3PuClG (melting point, 685 & 3") and a second compound (peritectic point, 611 f 3" a t the plutonium(II1) chloride mole fraction 0.35), possibly KzPuC16. The composition of the latter compound was chosen on the basis that it is the simplest compound consistent with the experimental data. The two eutectic points occur at the mole fractions 0.17 and 0.57 and the respective temperatures 621 f 3 and 489 f 3". An attempt was made to isolate each of the two compounds by crystallization from the appropriate liquid melts. An optical study of the samples obtained in this manner indicated the presence of two different phases distinct from pure potassium chloride and pure plutonium(II1) chloride. Acknowledgments.-We wish to thank W. J. Maraman and It. D. Baker of the Los Alamos Scientific Laboratory for discussions and encouraging interests. This work was done in the Los Alamos Scientific Laboratory. We are indebted to A. N. Morgan for the metallic plutonium, J. W. Anderson for the fabrication of the plutonium metal, C. F. Metz, G. R. Waterbury, C. T. Ape1 and L. A. Pulliam for chemical analyses and R. M. Douglas for inspection (optical and powder diffraction techniques) of crystals of the PuC13-KC1 compounds. T H E SOLUBILITY OF SILVER SULFATE IN ELECTROLYTE SOLUTIONS. PART 5 . SOLUBILITY I N MAGNESIUM SULFATE SOLUTIONS BY M. H. LIETZKEAND R. W. STOUGHTON COntribUtiOn from the Chemistry Division. Oak Ridge National Laboratory, Oak Ridge, Tenn. Received June 18, 1969
Previous papers in this series have described the solubility of AgzS04 in KN03,2 K2S04,*HPso44and "035 solutions. It was shown in these papers that expressions of the Debye-Huckel type could be used to describe the solubility data in each system over a wide range of temperature and ionic strength. Since it seemed to be a logical extension of this solubility program to determine whether similar expressions could be used to describe the solubilit,y of Ag2S04 in polyvalent electrolyte solutions, a study has been made of the solubility of Ag2SO4in 0.1, 0.5 and 1.0 m MgS04 solutions to above 150". Again a high speed digital computer has been used in making the calculations. (1) This paper is based upon work performed for the United States Atomic Energy Commission a t the Oak Ridge National Laboratory operated by Union Carbide Corporation. (2) M. H. Lietzke and R. W. Stoughton, THIBJOURNAL,63, 1188 (1959). (3) M. H. Lietzke and R. W. Stoughton, ibid., 63, 1186 (1959). (4) M. H.Lietzke and R. W. Stoughton, iMd., 63, 1188 (1959). (5) M. H.Lietzke and R. W. Stoughton, ibid., 63, 1190 (1959).
NOTES
Nov., 1959
1985
Experimental The solubility measurements were carried out using the same technique described previously.8 As in the case of the KNOt and KzS04systems, i t was fairly difficult to get a good reproducible set of data, since the solubility points were rather sluggish. The measurements were extended at each concentration of MgSOl to as high a temperature as possible. The upper limit was determined in each case by interference due to hydrolysis of both the AgZSO? and the MgSO,. It is believed, however, that the solubility data reported represent a t least a t the higher temperatures metastable solubility values for Ag,SOr in MgS04 solutions, since the establishment of solubility equilibrium was more r?pid than the onset of hydrolysis in the case of the points given. The measurements were reproducible to about & 2 O in 0.1 m MgSO,, and to about f 5 O in 0.5 and 1.0 m MgSOd solutions.
Results and Discussion I n Fig. 1 the circled points represent the experimentally observed solubilities of AgzS04 in H20 and in O.P,0.5 and 1.0 m MgSO4 solutions. The results in H,O up to 100" were obtained from the work of Barre,? while the values a t the higher temperatures were reported in the first paper in this series.2 The experimental solubility values of Ag2S04 in the MgS04 solutions are given also in Table I. TABLE I THESOLUBILITY OF Ag2S04IN MgS04 SOLUTIONS MgSOd, m
0.1
t
Bobsd
25 64 85 104 154 160 175 64 85 129 147 150 64 05 103 124 144
0.0247 ,0351) ,0422 .0441 .0506 .0512 ,0506 0.0440 ,0520 ,0664 ,0732 .0801 0.0520 ,0656 ,0718 ,0874 ,0990
~ C B I C ~
0.023 ( A s = 0.82) ,036 ,042 ,045 .051 ,052 .052 0.043 ( A s = 0.77) ,052 ,069 .076 .079 0.051 (As = 0.72) ,069 ,074
I 0
I
0.02
I 0.04
,
' 0.06
,
~ 0.08
,
/
0.10
, 0.42
l
, 0.f4
I 0.16
meq2s04.
Fig. 1.-The
solubility of AgZSOd in MgSOh solutions.
Ag2SO4 in pure HzO, equals 4sO3 a t ionic strength = 3s0; ST is the appropriate Debye-Huckel limiting slope a t the given temperature; and As is a concentration dependent but temperature independent parameter. I n computing the change of ST with temperature the equation for the variation of the dielectric constant D of water with temperature given by Akerlijf and Oshrys was used. At any temperature and concentration of MgS04 the solubility product X of the AgzS04 is given by
Io
s = 4sys + m )
(3)
In starting the calculations at each concentration of MgSOc the observed solubility of AgzSO4 a t each temperature was used to compute a value of I. Then using equations 2 and 3 a value of &lcd (the calculated solubility of Ag2S04) was ob0.5 tained. This value of Scaled was used to correct I (equation 1) and the process repeated until successive values of Scaled agreed to within 0.1%. The calculations were carried out a t 25" intervals from 25 to 200" a t each concentration of MgS04 using a 1.0 range of values of As. Then from a plot of Scaled 21s. As a t each temperature and concentration of MgS04 it was possible to find the value of As .087 which gave closest agreement with the observed .IO1 solubility. It was found that the value of As On the assumption of complete dissociation of varied little with temperature a t each concenboth electrolytes, the stoichiometric ionic strength tration of MgSOd but did show a decrease with inI of the AgzS04-MgS04 solutions is given at any creasing concentration of MgS04. In Fig. 1 the molality m of MgS04by solid lines indicate the calculated solubilities corresponding to the values of As averaged a t each I = 4m + 35 (1) concentration of MgSOd. The As values for each where s represents the molal solubility of AgZSO4 concentration of MgSO4 are: for m = 0.1, As = in the MgS04 solution. As in the previous papers 0.82; form = 0.5, As = 0.77; for m = 1.0, As = in this series the stoichiometric solubility product 0.72. The decrease in the values of As with coilof Ag2S04 on a molality basis X a t any molality m centration (and hence of the ion size parameter of MgSO4 was assumed to be given in terms of the 12 which is contained in the As) is consistent with ail solubility product in pure H2O by a Debye-Huckel hypothesis involving smaller hydration spheres expression of the type for the ions as the concentration of MgS04 increases. 47 Calculations also were performed assuniing an In S = In So + ST average temperature and concentration independ[l+AsdT 1 + A s d 6 1 (2) I n t,his equation 80, the molal solubility product of ent value of As ( = 0.77). I n this case, however, the solubility values calculated at m = 0.1 were (6) M. H. Lietake and R. W. Stoughton, J. A m . Chem. Soc., 78, 3023 (1956). (8) G. C. ifkerlof and H. J. Oshry, J . A m . Chem. Sac., 72, 2844 dzl
(7)M. Barre, Ann. chim. et phys., [SI,2 4 , 211 (1911).
(1950).
COMMUNICATION TO THE EDITOR
1986
too high, while those calculated a t rn = 1.0 were too low. Hence it appears that the value of As must be considered concentration dependent. It is evident that a Debye-Huckel type expression for the variation in the molality solubility product of Ag2S04 with ionic strength in MgSOa
Vol. 63
solutions can be used to fit the solubility data over a wide range of concentration and temperature. Acknowledgment.-The authors wish to express their appreciation to Mrs. Laura Cain Meers for performing the experimental solubility measurements.
COMMUNICfATIONTO THE EDITOR ON T H E RADIATION-INDUCED POLYMERIZATION OF ISOBUTYLENE IN THE LIQUID PHASE
sir: Recent studies'J of the low temperature, radiation-induced polymerization of isobutylene have established quite conclusively that the chain propagation is not free radical, but ionic, and that the most likely initiating species is the (CH&C+ ion. Free radicals are formed in the radiolysis, however, as evidenced by the disappearance of diphenylpicrylhydrazyl and other scavengers. The natural question as to the nature of the simultaneous modes of formation of (CH&C+ and free radicals in irradiated isobutylene seems to be resolved most easily by consideration of the known reactiona of the isobutylene molecule-ion with isobutylene; that we need consider only reactions of the molecular ion of isobutylene, and not those arising from ion fragmentation processes, is supported by the results of recent calculations by S t e v e n ~ o n ,which ~ suggest that for the liquid-phase time interval between C4H8+formation and collision with a C4H8 molecule (10-l2 second or less) such fragmentation will be negligible. Tal'roze and Lyubimovaa showed that a reaction producing (CHs)&+ in isobutylene does indeed occur, viz. i-C4Hs+
+ GCIHS +C4Hs + C4H7 +
(1)
in which, on the basis of energetic considerations of (1) W.H.T. Davison, S. H. Pinner and R. Worrall, Chem. and Ind. 1274 (1957). (2) E. Collinson, F. S. Dainton and H. A. Gillia, THISJOURNAL,6S, 909 (1959). (3) V. L. Tal'roze and A. K. Lyubimova, Dokladzr Akad. Nauk, S.S.S.R., 86,909 (1952). (4) D. P. Stevenson, Radiation Research, 10, 610 (1959).
butyl ions,b C4Hs+ can hardly be other than tertiary, and C4H7 is most likely a methyl substituted allyl radical; this radical has been suggested by Collinson, Dainton and Gillis2 as being present in the radiolysis. Further, the very observation of (1) in a mass spectrometer3 indicates its bimolecular specific reaction rate to be of the order of cm.s-molecule-l-sec.-l, which means that it most likely can compete favorably with neutralization reactions. Thus, if (1) occurs without competition, every ionization event results in the formation of a polymerization initiator, (CHa)&+, and a free radical, C4H7, with an energy yield of G[(C&)aC+I = G(CIH~) l O O / W C r H s
I
(2)
where Wc4a is the energy required to form an ionpair in isobutylene. Wc4&is not known but, taking the measured value6 of WC~H, and assuming the ratios of W to ionization potential to be equal for ethylene and isobutylene, can be estimated to be 23.6 e.v. From (2), this leads to a calculated G(C*H,) of 4.2 which is in quite good agreement with the experimental G(-DPPH) of 3.7 found for the disappearance of diphenylpicrylhydrazyl. This agreement lends support to the postulate that the ion-molecule reaction, (1))known to occur in the gas phase, is of great importance in the liquid phase polymerization of isobutylene because it represents the mode of formation of the polymerization initiator. RESEARCH AND DEVELOPMENT DIVISION OIL & REFININQ COMPANY HUMBLE BAYTOWN, TEXAS RECEIVED OCTOBER9, 1959
F. W. LAMPE
(5) F. P.Lossing, P. Kebarle and J. B. Deaousa, paper presented at the Institute of Petroleum Hydrocarbon Researah Group and A. 8. T. M.Committee E-14 Joint Conference on Mass Spectrometry, London, September, 1958. (6) J. Weiss and W. Bernstein, Phys. Rev., 103, 1253 (1956).
I 4
4
