THE SOLUBILITY OF SILVER SULFATE IX ELECTROLYTE SOLUTIOSS. PART 7 . SOLUBILITY I S URASYL SULFATE SOLUTIOh-S1 B Y 31. H. LIETZKEkND 1%.\v.
STOGGHTOY
Contrzbutzon from the Chemzstry Divzszon, Oak Rzdge ;Vatzonal Laboratoi y, Oak Rzdge, Y'enneasce Received January $3, 1960
The solubilitv of Ag,S04 has been measured in 0.100 to 1 348 m UOPSO4 solutions as a function of temperaturr to about 200" The agreement bet15 een calculated and observed solubilities \%asgood when hydrolytic and complexing reactions of the uranyl ion were taken into account The calculated concentrations of all assumed species are presented as functions of UozSO4 concentration and temperatnre, and it is concluded t h a t the relative stability of the neutral species UO+O1 compared to LO?++and U0,(S04),-- increases with temperature. Both the enthalpy and entropy for the asioeiation of UOj-+ and bO4-- into the neutral species appear to attain large positive values at elevated temperatures, indicating that a large degradation of solution stiucture occurs as the ions associate
Previous papers in this series have described the solubility of Ag2S04in KNO3, K2S04,H2S04,HNOj and MgS04 solutions.* It mas shown in these papers that single parameter expressions of t'he Debye-Huckel type could be used to describe t'he variation of the Ag2SOI act.ivity coefficient over a wide range of t,emperature and ionic strength. In each of t.hese cases complete dissociation was assumed except for the species HS04- and HXOa. The equations for the variation of the acid coiistants Ki"of these species with temperature were obtained from Young,3 while the variation of the acid quot,ients Ki with ionic st'rength I a t ally teniperature were also assumed to be given by single parameter Debye-Huckel expressions
Experimental The solubility measurempnts were rarriecl out ivith the same technique described previously.; The measurements were extended a t each concentration of U 0 2 S O pto as high a temperature as possible, the limit depend~ngon the temperature of appearance of B red hydrolJsis product of uranium containing silver. The values a t 1.348 m (molality of) UOnS04 below 174" were taken from previous work.6
Results and Discussion In Fig. 1 the open circles represent the experimentally observed solubilities of i i g a S 0 4 in H2O and in 0.100, 0.409, 0.622, 1.060 and 1.348 m UOZSO, solutions. The results in H 2 0 to 100" were obtained from the work of Barre,7 while the values at higher temperatures were reported in the first paper in this series.* The data in Fig. 1 show that uranyl sulfate causes a large enhancement of the solubility ln K , = 111 K , O+ ~ , r d f / ( + l A , ~ Z ) (1) of -4g804, particularly at the higher temperatures; this enhancement is appreciably greater than that where ST is the appropriate Debye-Huckel limiting shown by b1gS04.2 For example, the observed slope a t any t,emprature, Ai is the siiigle parameter, solubilities a t 150" are about 0.11 niid 0.40 molal in arid K i o represents the thermodynamic constant 1 m MgSOd and U02S04,respectively (The solid (at I = 0;. 111 all w s e s best agreement, bettween ob- heavy lines bhow t h r solubilities cdculated when served ax1 c~:ilculated solubilities was ohtaiiietl taking into account coniplexillg aiid hydrolysis; when earh siiigle A i parameter \viis assumed to he see section 1b below). temperatiire iiidepeiideiit aiid to be either iotiic: 1. Solubility Calcula'iions (a). Assuming Comstrength independciit or to decrease slowly ivith plete Dissociation. -- If both lr( )?SO1 and A4g2S04 increasing ionic: strength. :ire :wunicd t o be complctely di-;-ocwtcd the only The preseiit paper deals with the solubility of equilibrium ir hich iietd he coiisitlrrctl is :&SO4 it1 aqueous I.702S04solutions, the st,udy .\g2so, = 2Ag+ + so4-t2) having htwi undrrtaken in order to compare this medium with JIgS04 aolutions. Since the assump- The \-:wiatioii of the (niolality) solubility product tion of caoniplcte dissoc*iation i i i the case of 3fgS0.i S = m.ig+-'7nsoc- u her? 1 from the \-al\iei i i pure' i i a gave good agreement, l~etweenohserved and ral- ter S o = 4sod (it here 6 0 is the riio1:~lsolubility iii pure d a t e d solubilities :md s i n e l;02S04is known4 not \iater) is giveii by the ecluatioii to he c~omplt~tely clissoc-iated at, moderate coii(~:iitrations, a, coriiparisoii x i s t,hought to he of interest,. The follo\r-iiig iiidependent assunipt,ioiis \vere mad(! in attempting to explain t,he data: (1) cwniplcto dissociation of V02S04, ( 2 ) complesing of P O 2 + + SI is the Deb)-e-Huckel limiting ilope for ai1 ioii of with SO,---. (8) hydrolysis of IT02+i, aiid (4)both uiiit charge, aiid A , is the single paranwter. The value of the A , parameter in eq. 3 which gavr ~.omplexingaiid hydrolysis of KO2++. the best fit with the data is lower than any of those (1) This p;.per is based upon work performed for t h e United States for the other electrolytes studies in this series; Atomic Energy ('ommission a t the Oak Ridge S a t i o n a l Laboratory e.g., it is about 0.4 compared to about 0 7.5 for o i ~ e r a t e dby 1-nion Carbide Coyparation. hIgS04 solutions and 0.65 to 1.1 for the others. If ( 2 ) 31. €1, I.ietzke a n d R. W ,Stoughton, T I I I S.JoL-R~;.\L. 63, 1183, 118(i. 1188. 1190. 198L ( 1 0 3 9 ) . the As parameter is optimized at each T702S04 coii1 3 ) T. F. Young. prix-ate comnlunication: also T. F. 1-oung, 1.. 5'. centration it shows an increase with the latter from .\Iarsnville and H. .\I. S m i t h , C h a p t e r 4, "The S t r u c t u r e of Electrolytic ~~
Soliitions." edited by !IJ ..Hamer, J o h n Wiley and Sons, Inc., Kew 1-ork, N. Y . , 1!?.5$l, p p . ad-ii?. (41 R. A. Kraiis a n d F. Nelson, Ciiap. 23, " T i i ~Striictiirc o f P : l c ~ trolytic Sohitions." F;ditrri hy \V. J . H a m e r , John !Tiley a n d Sons, Inc., S e w Yurk, S . 1.. l ! l X , 1,. 349.
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June, 1960
SILVERSCLVATE ISUKANYL SULVATE SOLUTIOSS
~ O L U B I L I T YOP
0.23 a t 0 1 in to 0 43 a t 1.348 m I n all of the other media studied the optimized value of A , showed a decrease or no change with increasing supporting electrolyte concentration. The assumption of a temperature-independent, concentration-dependeiit A , shoved poorer agreeemeiit betn een observed and calculated wluhilities in the case of U02S04 than iii the i)ther cases The (act I ity) solubility product of silver sulfate K," waz evaluated at each 25" interval to 125" by evtrapolati ig to zero ionic strength as described previouslgy and compared n ith the average ralue obtained in t h e other media M7hile the maximum x d u e s in I
