M. H. LIETZKE AND R. W. STOUGHTON
4520
40000b
*
IO KNb03
Fig. 3.-Variation
'
20
'
'
'
'
30 40 50 60 MOLE PERCENT. KTa 0,.
'
70
' 80
'
I
90 100 KTa 0 ,
of lattice constant of KSb03 with mole % KTaOs a t 450".
lum concentration. This decrease in lattice constant can only be attributed to an increase in the ionic character of the heavy metal atom-oxygen bond. Using the data of Shirane, et aI.,26for cubic KNb03 a t various temperatures above the Curie point, one obtains a straight line relationship between cubic lattice parameter and temperature. The extrapolated value obtained for KNb03 a t (26) G. Shirane, R. h'ewnham
and R . Pepinsky, Phys. Rev., 96,
Vol. 78
room temperature as 4.0033 A. As the KTa03 parameter is 3.9885 A. this further substantiates the belief that the Ta-0 bond is more ionic than the Nb-0 bond. A qualitative application of the Huggins relat i o n ~ h i pindicates ~~ that "effective" ionic characters of the Ta-0 and Nb-0 bonds are essentially the same. The equation was applied, assuming that the contribution of K would be the same for both systems. What is most apparent, however, is that the combination of geometric and electronic factors is not significantly different for T a and Nb indicating that the Goldschmidt theory2s is extremely sensitive to this combination. Therefore one would rarely expect to encounter model behavior relationships. Acknowledgment.-The authors wish to acknowledge the invaluable assistance given by S. Triebwasser during his many helpful discussions and to thank Miss S. Silverman for preparation of samples, and the h'orth American Philips Company for making available equipment for several experiments. (27) M. L. Huggins, THISJOWRSAL, 76, 4123, 4126 (1953). (28) J. M. Goldschmidt, Skrtft Norskc, V i d . A k a d . Oslo. M a t . N u l .
KI., S o . 8,7 (1926).
NEWYORP, NEW YORK
581 (1954).
[CONTRIBUTION FROM
OAK RIDGENATIONAL LABORATORY]
The Thermodynamics of Indium Sulfate Solutions1 BY M. H. LIETZKE AND K.W. STOUGHTON RECEIVED MARCH26, 1956 E.m.f. data are presented for the cell In, In2(SO4)a(m), Ag&304(sat), Ag for Int(S0Ja concentrations of 0.01 to 2.3 m and The log of the activity of In2(S04)3 is approximately proportional to the molality a t any given temperatures of 15 t o 70 temperature. E.m.f., conductivity and solubility data all strongly indicate considerable ionic association in Inz(S0Ja solutions.
.
Introduction (0.02676 m in water a t 25') combined with the The cell In, Inz(S04)3(m),Hg2S04, Hg was studied necessity of keeping i t out of the In, In2(S04)3 by Hattox and DeVries2 as a function of I~:!(SOI)~electrode compartment imposed a further restricconcentration a t 15, 25 and 35'. They calculated tion on the lower limit of useful Inz(SO4)3 concenactivity coefficients of In2(S04)3on the assumption trations as far as interpreting the data was concerned, because of the problem of liquid junction of complete dissociation into I n + + + and SOI-ions a t the lower concentrations and by making the potentials. usual extrapolation using the Debye-Hiickel equaExperimental tion. The current study began as an attempt to The AgzSO4 used in all the measurements was prepared extend the useful range of the In, Inz(SOa)a(m) according to the method described by A r ~ h i b a l d while ,~ the electrode to higher temperatures. Because of the Int(SO& was prepared by the method suggested by Hattox lesser tendency of AgZS04 than Hg?S04 to hydro- and DeVries.2 E.m.f. Measurements.-The Ag, Ag2S04 electrode conlyze as the temperature is raised and the acidity sisted of a pure silver wire dipping into the desired In2(S04)s lowered and the consequently greater reliability of solution saturated with Ag2SOI. The indium electrode the Ag, AgZSOd electrode under these condition^,^ consisted of a platinum spiral plated with indium metal as the cell In, Inz(S04)3(m),AgsSOd(sat), Ag was used. previously described.* Both electrodes were contained in The In2(S04)5concentration was varied from about 10 mm. o.d. glass tubes which were drawn down to capillary at the lower end and then bent up in the form of a U. 0.01 to 2.3 m and the temperature from 15 to 70'. size The Inz(S0,)s solution and electrodes were contained in a The lower limit in concentration and the upper Pyrex tube fitted with a Teflon cover having holes to admit limit in temperature were fixed by the ability to ob- the electrode compartments. The entire electrode assembly was sealed in a steel bomb fitted with electrically insulated tain reproducible results. The relatively highly solubility of AgZS04 leads. As shown previously,s this type of electrode as(1) This document is based on work performed for t h e Atomic linergy Commission a t t h e Oak Ridge Natlonal Laboratory. ( 2 ) E. M . Hattox a n d T. DeVries, THISJOURNAL, 68, 2126 (1936). (3) M.H.Lietzke a n d R. W. Stoughton, { b i d . , 76, 6226 (1953).
(4) E. H . Archibald, "The Preparation of Pure Inorganic Substances," John Wiley and Sons, Inc., New York, N. Y., 1932. ( 5 ) M. H. Lietzke a n d J. V. Vaughen, THISJOURNAL,7 7 , 876
(1955).
Sept. 20, 1956
THERMODYNAMICS OF INDIUM SULFATE SOLUTIONS
sembly is very easy t o remove from the bomb for the purpose of changing solution or electrodes. T h e bomb containing the electrode assembly was immersed in a constant temperature bath controlled to a t least i ~ 0 . 0 5 " . The e.m.f. values were measured with a Rubicon precision potentiometer which was standardized with an Eppley cell calibrated by the Bureau of Standards. T h e time required for the cells t o reach equilibrium varied with the temperature. At the lowest temperatures the cells were allowed t o stand overnight although the potentials became steady after 3 to 4 hours. I n the intermediate temperature range equilibrium was established in 2 to 3 hours, and a t the higher temperatures within 2 hours. In all cases the e.m.f. measurements were carried to as high a temperature as possible. I n the most dilute solutions this was 40°, while in the more concentrated solutions 70". Above these temperatures irreversible hydrolysis of the Inz( S O C )occurred. ~ This was indicated by a sudden drop in electrode potential after which the system never recovered a t lower temperatures. I t was found that the Inz( solutions changed slowly with time, the rate of change being a function of the Inn( S O C )concentration. ~ Below 0.1 m it was necessnry to complete the e.m.f. measurements within one day. In the range 0.1 t o 0.3 m , the measurements could be extended to two days, while a t 0.5 m and above the measurements could be reproduced for a week. This effect also may have been due to irreversible hydrolysis of the Inn( solution. Solubility Measurements.-Aliquots of Inz(So4)~solutions of known concentration were saturated with AgZS04 by stirring the solutions with excess Ag,SOa crystals a t 25' on a magnetic stirrer. The solutions were then analyzed for silver volumetrically by means of a Dow Precision Recordomatic titrimeter. In carrying out the determinations, sample aliquots were added to a n excess of HCl and the excess chloride titrated with AgN03. It was found that this method of silver analysis gave much more precise and reproducible results than a gravimetric silver determination. The densities of the solutions were measured by weighing aliquots delivered from a Normax pipet. Conductivity Measurements.-The resistances of aliquots of the same solutions used in the AgZSOd solubility measurements (both before and after saturation with AgQSOa) were measured by means of a Serfass Conductivity Bridge. A dipping conductivity cell with a cell constant of 1.08 was used. Both the solubility and conductivity measurements were carried out a t 25" only.
4521
TABLE I THEEFFECTOF TEMPERATURE AND Ins(SOa)a CONCENTRATION O N THE e.m.f. (IN VOLTS) OF THE In, hz(S04)a(m;, Ag2S04(sat), Ag ELECTRODE SYSTEM 0.009888m E 1 3 . 2 1,1048 2 0 . 4 1.0977 2 7 . 5 1.0905 3 6 . 1 1.0837
t,OC.
0.1990 m t,'C. E 1 5 . 0 1.0923 2 0 . 0 1.0880 2 5 . 1 1.0836
0.04933 m t,T. E 1 4 . 5 1,0992 2 3 . 9 1,0901 2 8 . 3 1.0862 3 5 . 6 1,0795
0.3959 m 0.5527 m E t,"C. E 1 3 . 3 1.0915 14.6 1.0830 20.2 1,0838 3 3 . 2 1.0661 3 0 . 0 1.0747 4 2 . 4 5 1.0577 3 5 . 3 1.0701 5 0 . 5 1,0502 4 1 . 8 1.0644 58.3 1,0434 4 9 . 9 1.0576 68.9 1,0342 5 7 . 3 1.0518 7 4 . 6 1.0291 6 5 . 4 1.0450 7 0 . 5 1.0409 1.2909 m
1,OC.
0.9732 m E 1.0599 1.0422 1.0310
I, o c .
24.9 48.2 63.4
0.09799 m I,'C. E 14.3 1.0960 2 4 . 4 1.0870 4 0 . 3 1.0730 5 0 . 2 1.0647
t,
oc.
14.0 19.3 29.5 40.8 50.2 60.3 70.5
0.1957 m
E
1 . T .
2 4 . 9 1.0841 3 6 . 2 1.0735 4 4 . 8 1.0661 5 0 . 4 1.0616 6 0 . 3 1.0533 0.6806 m f,OC. E 1 3 . 6 1,0791 2 2 . 7 1.0706 2 9 . 8 1.0645 3 0 . 4 1.0640 3 9 . 5 1,0558 5 0 . 6 1,0468 6 2 . 3 1.0370 7 1 . 5 1,0299 2.2822 m
E
I, o c .
E
1.0588 1.0531 1.0449 1.0368 1.0306 1,0245 1.0176
12.5 21.8 23.1 29.0 39.4 40.2 49.8 70.5
1.0296 1.0225 1.0215 1.0175 1.0105 1,0102 1,0053 0.9951
Vries are more positive. The discrepancy is about 5 millivolts a t 0.4 m and 15' and becomes increasingly larger a t lower concentrations and higher temperatures, amounting to about 20 millivolts a t 0.01 m and 15' and to about 8 millivolts a t 0.4 and 35'. Thus the e.m.f. values measured by Hattox and DeVries appear to be somewhat too high because of hydrolysis of the Hg, HgZS04 reference electrode. Since the Ag, Ag2S04 electrode was shown t o be less susceptible to hydrolysis than the Hg, HgzResults and Discussion E.m.f. Measurements.-Table I summarizes the Sod, i t seemed that the Ag, AgZS04 electrode should experimental measurements of the e.m.f. of the be a better choice as a reference electrode and that cell In, Inz(S04)a(m),Ag2SO4(sat), Ag as a func- the calculations involving the thermodynamic functions of Inz(S04)3solutions should be more nearly tion of In2(S04)3concentration and temperature. In previous measurements of the thermodynamic correct in the present work. In both cases the reproperties of In?(S04)3 solutions,2 a saturated Hg, sults may be slightly in error due to hydrolysis of ) ~ would lower the In2(S04)3 HgzS04 reference electrode was used. However, the 1 n ~ ( S 0 ~which HgzS04 is known t o hydrolyze in weakly acid sys- activity by the amount hydrolyzed. The ratio temss and a t higher temperatures. For example, moles of acid formed per mole of Indium (H+)/ when the cell Hg, HgZSOd(sat), HZSO,(m), AgzS04- (Inf3) increases a t 23' from about 0.035 a t 0.2454 (sat), Ag was studied,s i t was found difficult to m to about 0.12 a t 0.0063 m.2 Since the extent of attain the theoretical e.m.f. for the cell a t 25' in hydrolysis is relatively small and since attempts to acid solutions more dilute than 0.05 m. This ef- reduce the hydrolysis by the addition of acid were fect which was even more pronounced a t higher unsuccessful in producing consistent e.m.f. data, temperatures, was shown previously to be due to no attempt was made to correct for hydrolysis in formation of HgOS3 The discrepancy is in such a the subsequent calculations. Added acid dissolved direction as to make the e.m.f. of the In, Inz(SO4l3- the In metal of the In electrode so rapidly that no (m), Hg2SO4(sat),Hg cell too positive if any hy- stable potential was obtained. In the course of the work no attempt was made drolysis of the Hg2S04 occurs. to obtain e.m.f. measurements a t even temperaWhen the e.m.f. values measured in the present work are corrected for the difference between the tures. For convenience in calculating and presentAg, AgZS04 and Hg, HglS04 electrode potentials ing data on activity coefficients and partial molal and compared a t the same concentration with the heat contents the e.m.f. values were plotted against values obtained by Hattox and DeVries i t is found temperature a t each experimental concentration to that the e.m.f. values reported by Hattox and De- obtain values a t 5' intervals. The e.m.f.'s a t the even temperatures were also plotted as a function of (6) J. W. Mellor, "A Comprehensive Treatise on Inorganic and Inz(S04)3concentration to obtain values a t even Theoretical Chemistry." Vol IV, Longmans, Green and Co., London, concentrations. 1923. p. 966
h1. H . LIBTZKE AND R.IV.STOVCI-ITO~-
4.525
II'hen the In2(S04)3 electrode is measured against the saturated .4g, Xg,S04 electrode the cell reactioti is 2111 + :jAg&OA I_ 11i?(SOcI , ~ ( + ~ ~ IiAg ) (I The e.ni.f. is given by the xernst equation i n the form
E
; (2) trations a t least, although the evidence is not neceswhere u I n 2 ( ~ ~represents ,)a the thermodynamic ac- sarily definitive. Hence, as a consistency check on tivity of the In2(S04)3solution. In order to coni- the e.1n.f. data, the left side of equation 5 was pute activity coefficients from the e.m.f. data it is plotted against the square root of the Inr(S04):1 molality to give the :7 7 family of curves shown in Fig. 1 . Because of the curvature of the lines a t lower conI centrations this type of plot is not useful for extrapolation to obtain E" x-alues for the electrode combination. H o w ev er , t h c curves do indicate the consistency of the e.ni.f. data and show a t each temperature the range o f concentration in which the activity coefficients decrease in value then in0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 1 . 1 1.2 1.3 1.4 1.5 1.6 crease again. v;;;In an effort to the Fig. 1.-Plot of E RT/3F In nz 21s. d m for the In, Inl(S04)3(m'i, Ag2SO4, Ag electrode combination. data to obtain Eo values for the elecnecessary to make an assumption concerning the trode combination and hence to permit the calspecies in solution. Three different assumptions culation of activity coefficients for the In2(S04):I which might be made are that Inz(SO?)?dissociates solutions equations 3, 4 and 5 were modified by substitution of the Debye-Huckel expression for completely into 2 I n + - + + 3S04-(a) the In y term and rearranged to give a t 5.5' the excompletely into 21nSOc- + SO4-(b) completely into InS04 + In( (c) pression
,
+
+
Depending on the assumption made the activity of h ( S O d 3 would be given by (a) 108m5y6, (b) 4 m 3 y 3 or (c) m 2 y 2respectively, , where y is the stoichiometric mean ionic activity coefficient consistent with the particular assumption made. Of course, the actual species existing in solution may not correspond to any one of the above assumed modes of dissociation, or even to a combination of them; that is, the actual predominant species may be more complex than those indicated. The hope was that by attempting to calculate the activity coefficients on one or all of the three bases mentioned, EO Values might be obtained and by comparing the 7"s with those of strong electrolytes conclusions might be drawn as to the species existing in solution. On the basis Of assumption (a), (b) or (e), eWation 2 might be rewritten a s equation 3 , 4>or 5 , after rearrangement of terms. E
T 5RT + RHF -- In 108 + In nz C,F
=
5RTln
E" - ~"( (3) RF
E""
=
[
E + o,()()F)x,~cJ log k
+
O..?;i999 1
log
IIZ
-
(T
v,z;') 32'
4 E
+ 1.nfj C v l z ,
For a 1-1electrolyte k
(
=
1, ZJ
z/nl
j
=
2 , and
]
(6) =
2;
vjzi2 =
(i;
d
j
2
j
for a 1-2 electrolyte k = 4,
V =
3, and i
= 108, y = 5 , and while for a 3-2 electrolyte v J j z = 30. U7hen the values of the constants j
were substituted into equation 6 for each case the curves shown in Fig. 2 were obtained. As can be seen subtraction of the Debye-Hiickel term does not in any case give a function providing a linear extrapolation to infinite dilution. Hence values of
Sept. 20, 1956
Eo for the electrode system could not be obtained by this method. Because it was not possible to derive values of Eo by extrapolation, activities and activity coefficients had to be expressed as ratios and partial molal heat contents as differences. For reasons discussed in the section on conductivity measurements all ratios and differences were referred to the value a t 0.1 m. Figure 3 shows the variation of log U / U ~ for . ~ I~I~(SO solutions ~ ) ~ as a function of concentration . ~ and temperature. The values of log U / U ~ were calculated by means of the expression log ala0
1
=
6F 2.303RT ( E o
~
3.0200
0 .
i'
(7)
- E)
=
?/?(Q.I)
WT 6F ( E ( " - E ) - log
' 0400L4 I
4.0300
which follows immediately from equation 2. The ratios are much more sensitive to temperature variation a t the higher concentrations than a t the lower. At each temperature, however, the log of the activity ratio is approximately proportional to the molality, a fact difficult to explain simply. Activity coefficient ratios y / y(o.l)were calculated for I n ~ ( s 0 on ~ )the ~ basis of a 1-1, a 1-2, and a 3-2 electrolyte a t 25' using the general equation log
4523
THERMODYNAXICS OF INDIUM SULFATE SOLUTIONS
1.0400
0 9800
-/
0 9700
-
0 9600'
L..
04
0
m
(8)
1)
l
l
l
l
l
'
~I
03
of E"" vs. m for the In, In2(SO&(m), AgzSOd, Ag electrode combination.
Fig 2 -Plot
The values of y 'y(o1) as a function of m are given in
i
02 m
'!
l
l
Il
)
l
/
I
j
]
(
]
1
1
0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 1.1 1.2 1.3 1.4 1.5 1.6 1.7 1.8 1.9 2.0 2.1 2.2 2.3 Fig. 3.-The
m. variation of log ala0.1for In?(SO& solutions as a function of concentration and temperature.
Table 11. The values of y/y(o.l) for the case of the 1-1 electrolyte increase more rapidly with m a t the higher concentrations than do those of any known 1-1electrolyte. This may be considered as evidence against any conclusion that In~(S04)3behaves like a strong 1-1 electrolyte throughout any large concentration ranse. Values of LZ - &.I) were calculated directly
from Fig. 3 showing the variation of log U / U ( O J ,with concentration and temperature. Since (d logdapo.")
=
- E% ___ - ZZ(O.1)
m
-
L,
-
Lz(0.1) = -2.303RT*
2.303RP
(9)
* logAT
(10)
Table I11 shows the temperature and concentration
nr. H. LIETZKEAND R. LV. STOUGHTON
4524
Vol. 75
the bases of dissociation as 1-1, 1-2 and 3-2 electrolytes). These curves are plotted in Fig. 4 along with the curves* of other electrolytes TIIE for comparison. Of the latter none except strong Y /? 0.1) m 1.1 1-2 3-2 acids and alkalies shows equivalent conductances 0,4210 0.3 0.5976 0.4919 greater than about 200 cm.2 ohm-' equiv.-' a t the ,3357 .i) 0.7310 ,4745 lower concentrations. None of them decreases as .7 1.224 ,5983 ,3374 rapidly with increasing concentration as does the ,3629 1.0 2.509 ,8570 Inz(S04)3 conductivity if this electrolyte is assumed ,4154 1.2 4 . G25 1.213 to dissociate as a 1-1 or 1-2 electrolyte. Also none ,5303 1.5 11.89 2.113 shows as low conductivity as does on the ,7755 2.0 47.3; 4.786 basis of complete dissociation as a 3-2 salt. 2.2 77.22 6.471 ,8904 If In2(SO& were assumed to dissociate into about 2.3 123.5 8.718 1.046 three ions (e.g., 2InS04+ and sod--) a t 0.01 m and dependence of 1 2 - ~ Z C O . ~ , values calculated by to dissociate less and less with increasing concenmeans of equation 10, from the table it is evident centration, then its equivalent conductance based that the 1 2 - 12(0.~) values become more tempera- on the assumed ion concentrations would show a ture dependent a t higher concentrations of In2- behavior qualitatively similar to that of the other salts. The conductivity data are therefore inter(soda. preted as evidence against complete dissociation of TABLE I11 I n ~ ( S 0 4into ) ~ I n + + + and so,-- ions in the concentration range studied. These data also imply a VALUES OF 1 2 - Z,,,.,, IN KCAL. FOR I n ~ ( S O SOLUTIONS ~)~ AS A FUNCTION OF CONCENTRATION A N D TEMPERATUREdecreasing degree of dissociation with increasing najt 200 30' 40° 50' concentration, as might be expected. TABLE I1 I'ALUES FOR In2(S04)3 SOLUTIONS CALCULATED OK BASIS O F A 1-1, A 1-2 AND A 3-2 ~ L B C T R O L Y T RA T 25'
r / y c o ,).
0.5 1. 0 1.5 2.0 2.2
2.0 6.3 9.0 14.2 14.9
2.5 7.6 12.0 13.9 15.1
2.7 8.1 13.5 19.3 20.6
2.9 8.1 14.3 19.1 20.6
THECONDUCTIVITIES m
TABLE IV O F Inz(SO4)sSOLUTIONS AT 25' Conductivities (ohm-' cm.-l) Without With AgSOc AgtSOl(sat.)
0.01 0.00242 0.00600 Conductivities of In2(S04)3 and Inz(S04)3-Ag2S04 .03 ,00527 ,00777 Solutions.-The conductivities of Tn2(S04)3solu.I ,0128 ,0159 tions both saturated with Ag2S04 and without .3 ,0267 ,0282 Ag2S04 were measured in order to get some idea of 1 . 0 .0347 ,0340 the conditions under which liquid junction potentials might be expected to be appreciable' and in Solubility of AgZSO4in Solutions.-The order to elucidate the degree of dissociation of solubility of AgzS04 as a function of conIn2(S04)3. Table IV shows the conductivities centration a t 25" is shown in the first two columns of the solutions as a function of Inz(SOa)3concen- of Table V. The solubility shows an increase with tration. It is seen that below about 0.1 m 1x12increasing In2(so4)3 concentration rather than a (SO& the conductivities of the solutions vary decrease as expected from the common ion effect. considerably depending on whether or not AgzS04 TABLE V is present (at saturation). I n this range i t was felt that the liquid junction potential between two THESOLUBILITY OF 4g2S04IN I I ~ ~ ( S SOLUTIONS O~)~ A r 25" Exptl. solutions of Inz(S04)3, one containing saturated Calcd. Solubility solubility Basis Basis Basis Ag2S04and the other not, may well become apminl(sor)t mAsdOI (b) (C) (a) preciable; and even the change in this potential ... ... 0.0263 0.02676" 0 with indium concentration may become comparable 0,0248 0.0252 (0.02675) ,02675 0.0096 in magnitude with the change in observed e.m.f. ,0222 ,0231 ,0277 ,0294 ,0272 Hence, any interpretation of the e.m.f. data in this ,0207 ,0240 ,0287 ,0494 ,0285 range may be meaningless. At higher concentra,0178 ,0196 ,0305 ,0989 ,0294 tions of Inz(S04)3 the conductivities depend very .0134 ,0157 ,0353 ,3020 .0341 little on the presence or absence of AgtS04, and it ,0092 ,0118 .0428 ,9809 ,0373 has been assumed that liquid junction potentials a G .Akerlof and H. C. Thomas, THISJOURNAL, 56, could be neglected without serious error. At 0.1 m 593 (1934). Inz(S04)athe effective equivalent conductance of the AgZSO4(calculated from the difference in con- Since the effect of increasing the ionic strength is to ductivity with the silver present and absent) was increase the solubility, the net effect of increasing about what might be expected on the assumption the ionic strength and the common ion concentrathat neither salt greatly influenced the ionic con- tion was estimated on the two bases of In2(S04)3 dissociation previously given as (a) and (b). The centrations resulting from the other. The "equivalent" conductances of pure Inz- AgZS04 was assumed to be completely dissociated (so4)3 solutions were plotted against the square root into silver and sulfate ions under all conditions. of the molar concentration on the bases of one, The solubility product was calculated on a molality two and six equivalents per mole (;.e., on basis for the pure Ag2S04solution and the change of (7) See G. N. Lewis and M . Randall, "Thermodynamics," McGrawHill Book Co., New York, N. Y., 1913, pp. 388-99.
( 8 ) B. W. Conway, "Electrochemical New York, N. Y., 1953, p. 142.
Data," Elsevier Publ. Co..
THERMODYNAMICS OF INDIUM
Sept. 20, 1956
380
4
452e5
SULFATE SOLUTIONS
k
t-
2201 \ 200
140 4 20
400
,-
O'
0'4
0'2
0!3
0!4
0'5
-
0'6
017
0'8
0'9
1'0
CZ
Fig. 4.-The
"equivalent conductances" of In*(S04)a solutions.
this product with ionic strength on bases (a) and (b) was determined using the modified DebyeHuckel equation for the log of the activity coefficient of an ion y : as shown in equation 11
with increasing In2(S04)3concentration. Hence, i t may be concluded from the solubility results that In2(SO& under the conditions reported is dissociated into fewer sod-- ions than indicated by basis (b). Although equation 11 is only approximate, most electrolytes known to be completely In this equation Zi is the charge of the ion in ques- dissociated deviate a t the higher ionic strength tion and I is the ionic strength. The constant 1.5 values in such a manner as to make the activity in the denominator was suggested to the authors by coefficients larger than the calculated values. ConScatchardg; the equation in this form gives a rough sequently, we do not feel that deviations from equaaverage value of the activity coefficient of an ion of tion 11 which would invalidate our conclusion are a strong electrolyte out to about an ionic strength likely to occur. It was then assumed that Inz(S04)3, a t least a t of unity, as can be verified by a plot of (l/Z+Z-) the higher concentrations, dissociates only into log y =tV S . Illafor various strong electrolytes. Using the molality solubility products calculated complex sulfate ions and hence causes no common in this manner, solubilities were calculated on bases ion effect. Using basis (c), equation 11, and the (a) and (b) and these are presented in the third and experimental solubility a t 0.0096 m In~(S04)3the fourth columns, respectively, of Table V. The Ag2S04 solubilities in the other solutions were calsolubilities calculated on both bases show a decrease culated; the calculated values are given in the fifth column of Table V. As can be seen, the agree(9) G. Scatchard, private communication
ment between the calculated and experimental .'lgzS04solubilities is surprisingly good. This may be interpreted as evidence in favor of the validity of basis ( c ) for the dissociation of In2(S04)3 a t the higher concentration, although actually one can only conclude that the magnitude of the ionic strength is consistent with basis (c). Summary and Conclusions The e.m.f. of the cell In, In2(S04)3(m),Ag2S04(sat), Ag was studied as a function of Inz(S04)3 concentration and temperature. Since it did not appear feasible to extrapolate to zero concentration to obtain Eo values relative activities only were obtained based on the value a t 0.1 m. The log of these ratios was approximately proportional to the concentration a t any one temperature. No unique interpretation in terms of solute species could be
[CONTRIBUTION FROM
THE
made. E.m.f., conductivity and solubility data all indicate that Inz(S04)3 is not completely dissociated into and sod= ions ill :my part o f thc concentration range studied. Thc conductivity and solubility data are qualitatively consistent with dissociation of In2(SOJ3 into InSO4+-and In(s04)~ions a t the higher concentrations and with further dissociation of these species in the more dilute range. On the other hand, a t the highest concentrations the rapid change of the activity coefficient of Inz(SOJ3on a 1-1electrolyte basis with M and the low magnitude of A,, appear to be inconsistent with expectations for a strong 1-1 electrolyte. Acknowledgment.-The authors wish to express their appreciation to Dr. T. F. Young of the University of Chicago for interesting discussions concerning this work. OAKRIDGE,TESVESSEE
CHEMISTRY DEPARTMEST, UNIVERSITY OF MICHIGAS]
Heat Capacity and Thermodynamic Functions of Uranyl Chloride from 6 to 350'K.I BY ELLIOTT GREENBERG AND EDGAR F. WESTRUM, JR. RECEIVED APRIL23, 1956 The heat capacity of pure anhydrous uranyl chloride was measured from 6 to 350'K. and the entropy, enthalpy and free energy function calculated from these data. At 298.16"K., the molal values of C, So and Ho- If: are 25.78 =k 0.04 cal./ deg., 35.98 =I=0.05 cal /deg., and 5157 & 8 cal., respectively. There was no detectable evidence of anomalous heat capacity behavior or significant magnetic contributions to the thermal properties.
Despite the increasing scientific and technological interest in uranium chemistry, the chemical thermodynamic data on important uranium compounds are in many instances inconsistent, estimated or unknown. Furthermore, the thermodynamic functions based on thermal data are frequently rendered insecure by the possibility of substantial magnetic contributions a t temperatures below those of the heat capacity measurements. The heat capacity of anhydrous, crystalline uranyl chloride (U02C12) was determined as part of a continuing study of the thermodynamic properties of the compounds of the actinide elements. The data and functions derived herein should prove of considerable utility ill chemical calculations when supplementary thermochemical data become available. Experimental Preparation of Uranium Tetrachloride.-Uranyl nitrate hexahydrate (500 g., Mallinckrodt A.R. grade, No. 8640) was treated with a large excess (850 ml.) of concentrated hydrochloric acid, and the mixture was evaporated t o dryness. The resulting uranyl chloride was refluxed for an hour a t 150' with four times ?he stoichiometric quantity (1 kg. ) of hexachl~ropropene.**~ The supernatant liquid was then drawn off through a sintered-glass filter stick, and the uranium tetrachloride was washed repeatedly with carbon tetrachloride until the washings were colorless. This procedure was carried out in a closed system in order to protect _~~______ (1) Adapted from a portion of t h e thesis of Elliott Greenberg, submitted t o the Graduate School of t h e University of Michigan in partial fulfillment of the requirements of the Ph.D. degree. This work was supported in part by t h e U. S. Atomic Energy Commission, Division of Research, Sub-contract A T ( l l - 1 ) - 7 0 , No. 5 . ( 2 ) Reported in: J. J. Katz and E . Rabinowitch, "The Chemistry of Uranium," McGraw-Hill Book Co., Inc., New York, N. Y., 1951, pp. 468-470. (3) J . J . Katz, private communication.
the uranium tetrachloride from atmospheric moisture, Most of the remaining carbon tetrachloride was removed by evacuation, and the last traces by heating to 250-300" under high vacuum for several hours. The anhydrous material was sublimed twice under high vacuum in a Pyres glass tube a t about 495'. Ignition in air of the sublimed dark-green crystals indicated 62.38 and 62.6270 uranium (theoretical 62.677,). Gravimetric analysis for chloride, by precipitation with silver nitrate, indicated 37.26 and 37.2770 (theoretical 37.3370). Preparation of Uranyl Chloride .-Finely ground uranium tetrachloride was converted to uranyl chloride (U02C12). in a system protected from moisture, by reaction a t 300350' with oxygen dried over anhydrous magnesium perchlorate. Attempts to minimize sintering by gradually increasing the temperature were unsuccessful; consequentlJ-, the procedure of grinding and treating with oxygen was twice repeated in order to effect complete conversion of the sample. The bright yellow product was analyzed by the methods already indicated, and found to contain 69.68 and 69.83% uranium (theoretical 69.827,), and 20.65 and 30.72% chlorine (theoretical 20.799&). Throughout the course of this work all transfers of hygroscopic material were performed in a drj7 box. Cryogenic Technique.-The Mark I adiabatic cryostat used for these measurements has been developed previously.4 The calorimeter (Laboratory Designation W-9) is similar in design and dimensions to W-65 except for two modifications: only four conduction vanes were used, and protection against possible corrosion was achieved by a 0.03 mm. gold plate on the interior surfaces. Apiezon T vacuum grease was used t o provide thermal conduction in the thermocouple sleeve and in the thermometer-heater well. The calorimeter contained 2.0 cm. helium pressure t o improve thermal conduction to the sample. A 150-ohm constantan, glass-fiber insulated wire was bifilarly wound in a conforming double thread in the heater sleeve. Temperatures were measured with a capsule-type platinum resistance thermometer (Laboratory Designation -4-3) (4) Edgar F. Westrum, Jr., and A. F. Beale, Jr., t o be puhliqhed ( 5 ) E. Benjamius and E. F. Westrum, Jr.. t o he puhlisheil
