GALENR. FRYSINGER S N D HENRY c. THOMAS
224
solute.1g Coefficients determined are included in Table I. Molar refractions for solids were calculated using the atomic refractions of Eisenlohr.20 Values for liquids were com(19) I. F. Halverstadt and W. D. Kumler, J . Am. Chcm. Soc., 64, 2988 (1942). (20) F. Eisenlohr, 2. p h y s i k . Chcm.. 76, 585 (1910).
1-01. 64
puted from the experimental densities and refractive indices.
Acknowledgment.-The authors are grateful to Research Corporation for a grant supporting this work.
ADSORPTION STUDIES ON CLAY MIXERALS. VII. YTTRIUM-CESIUM AND CERIUM (111)-CESI UM ON AIONTMORILLOSITE1 BY
GALENR. FRY SINGER^ AKD HENRYc. THO;M.~S~
Contribution N o . 1512 f r o m the Sterling Chemistry Laboratory, Yale University, New Haven, Conn. Received August 1.9, 1969
Exchange-adsorption of cesium and the nominally tripositive ions of yttrium and cerium on Chambers, Arizona, montmorillonite (from solutions of the nitrates) has been studied a t different concentrations and temperatures. The isotherms obtained with the two trivalent ions are closely similar but are very much more complicated in character than those previously found for ions of lower charge. Reversals in selectivity occur at compositions which depend on total concentration and temperature. The average charge of the sorbed trivalent ions is shown to be nearly three. Thermodynamic data for the exchange reactions are computed. The composition of the clay as regards the sorbed species is independent of the pH of the equilibrating solution over a wide range (3-4 pH units). rium column” method. In this procedure the adsorbent is brought to chemical equilibrium with a mived electrolyte solution of known composition. The solutions used in equilibrating the columns carried isotopic tracers; when the activity of the effluent became identical with that of the input solution, equilibrium was presumed to have been reached. Analysis of the column was done by isotopic evchange: the radioactive tracer was displaced by non-radioactive solution of the same chemical composition, and thus disturbances in the adsorption e uilibrium were avoided. The information obtained on the Iistribution of traced element between solution and adsorbent is independent of the nature of the molecular species present in the solution or on the adsorbent, provided only that these come to isotopic equilibrium. This last we may safely suppose to be true in our cases. The montmorillonite used was a portion of the material prepared and described by Gaines.s Determinations of loss in weight on ignition were done periodically in order to bring all weighings to a common basis. The columns contained about 2 g. of clay dispersed on 3 g. of an asbestos fiber shown to have a capacity of less than 0.002 meq./g. The cesium nitrate was stated by the supplier4to be 99.9 % pure. Solutions were prepared by direct weighing of the material dried a t 110’. Carrier-free C3137 from the Oak Ridge National Laboratory was obtained in HC1 solution. A volume of about 2 ml. containing 10 mc. was diluted to one liter and a few milliliters of this solution used as a “spike” for the cesium nitrate solutions, the chloride thus introduced being entirely negligible. The rare earth oxides from our sample of yttrium nitrat@ were found by radioactivation analysise to contain a miniThe techniques used in this work are essentially the same mum of 99.5 yo YzOJ, the impurities being rare earths and as those described by Gaines and Thomas,si.e., the “equilib- probably present only as traces. Solutions were prepared by dissolving the hydrated yttrium nitrate in very dilute nitric (1) T h e material for this paper is taken largely from the dissertation acid (pH 4-5). The resulting solution was passed through a submitted in 1956 b y Galen R. Frysinger to the Faculty of the Graducolumn of Dowex-1 in the OH form. By carefully controlling a t e School of Yale University in partial fulfillment of the requirements the flow rate it was ossible to obtain a stable solution of for the degree of Doctor of Philosophy and, as regards cerium, work yttrium nitrate of p g c a . 6.8 after bubbling with nitrogen. done b y Galen R. Frysinger in post-doctoral capacity a t Yale. The These solutions would stand fo’ weeks without showing eviwork was supported by the Nuclear Engineering Department of dences of precipitation. I t was not possible to accomplish Brookhaven National Laboratory, t o which we again offer our thanks. this with solutions of the chlorides, hence our choice of the (2) Venable Hall, University of North Carolina, Chapel Hill, North nitrate systems. Carolina. The Y9l isotope was used as a tracer. I t TYBS received in (3) G. L. Gaines and H . C. Thomas, J . Chem. Phys., 23, 2322 (1955). 1.5 N IICl and treated as follows: A solution containing 10 In this publication a factor of 2 was omitted in the definition of Kc’ me. of 1 - 9 1 was washed into a polystyrene beaker and a small on page 2325 as well &s in the calculations with this formula. All the portion of previously prepared Y( NOa)a solution added. values of Kc’ given in Table V must be multiplied b y 2 a n d the points The solution was made basic with NaOH and the resulting
In earlier papers of this series we have reported studies of the exchange-adsorption behavior of the alkali and alkaline earth ions on montmorillonite and of the alkalis on attapulgite. We now extend the survey t o tripositive ions. The results of all the earlier work are qualitatively similar. With cesium as the reference ion the adsorption behavior shows no particular anomalies and differs from system to system principally in the degree of selectivity for cesium a t low cesium content of the solution. The adsorption isotherms (except possibly for solutions containing only trace quanti: ties of one or the other of the elements) are all concave down over the whole composition range. With the tripositive ions this simple behavior disappears; the isotherms become much more complex in character. It is shown that the adsorbed trivalent species actually have a charge near three. The complex shape of the isotherms appears to be characteristic of the sorbing surface and little if a t all dependent on the ionic species in solution, since the composition of the surface is scarcely affected by changing the pH of the equilibrating solutions. Experimental
in Fig. 1 shifted down by 0.693. error: actually values of
-
so1
The heading in Table V I is also in
ln&’d(q/go)
are tabulated, and these
must be decreased b y 0.69. These changes give AFO = 3700 cal. and AS0 = 9 e.u.; A H 0 and the activity coefficients remain unchanged.
+
+
(4) A. D. Mackay, Inc., New York, N. Y., ( 5 ) From the Lindsay Chemical Co.. West Chicago, Ill. (6) We are indebted t o the Oak Ridge National Laboratory for this
analysis.
YTTRIUM-CESIUX AND CERIVM(III)-CESIUM ON MONTMORILLOXITE
Feb., 1960
225
I precipitate collected on a small circle of Whatman KO.40 filter paper. The filtrat contained negligible activity and was discarded. The precipitate was treated with 3-ml. portions of 0.1 N " 0 , until only a small amount of activity remained on the paper. The nitric acid solution was diluted to 500 ml. Stock solutions of yttrium nitrate were analyzed by the slow precipitation of the hydroxide. About 0.3 mmole of Y and 0.4 g. of urea in 200 ml. of water were heated on the steam-bath. After a few hours the precipitate appeared and and the di estion was continued overnight. The precipitate was collected on a Whatman KO.42 filter, washed with faintly basic IL",NOz solution, dried, ignited in platinum and weighed as YzOa. Solutions containing tracer were, of course, analyzed only after the preparation was completed. Cerium nitrate,5 stated to have a purity of 98% was dissolved and adjusted to p H near seven as described for yttrium. These solutions were analyzed by precipitation and repeated evaporation in platinum with HF. The residue was dried a t 300" and weighed as CeF3. The eluates of the columns were analyzed for cesium or for the trivalent metals by standard counting techniquep using glass jacketed Geiger counters, all appropriate corrections being applied. A sufficient number of counts was recorded to keep the statistical uncertainty in the neighborhood of 0.5 %. When solutions containing tracers for both elements were used in some of the experiments with yttrium, the yttrium was separated by a double precipitation as the hydroxide. The efficacy of this separation was demonstrated. The clay compositions were generally determined by the isotopic exchange Cso + Cs*. These exchanges gave relatively sharp break-through curves. The hold-up volumes of the columns were determined by weight. These volumes -0.100 amounted to 10-15 ml. 14 is impossible to state with what accuracy the pertinent, z.e., final, hold-up volume was known. Fortunately, however, errors of as much as 20% in 0.2 0.4 0.6 0.8 these volumes have little or no effect on the measured isotherms. The matter is of more significance in attempting to (44. detefmine the true charge number of the adsorbed trivalent Fig. isotherms, Cs-RE, on montspecies. We shall return to this point. 0, yttrium; b or 9 cerium? (a) 30°; The exchange capacity of the montmorillonite n-as p d e - morillonite at750. 0.02 termined for the nitrate systrm. Five isotopic cesium ex- (b) 500; changes and a Cs (natural ion) exchange gave 1.377 meq. may be seen. The eschallge capacities of the per g. backbone with a maximum difference from this average of 0.014 Gaines3 measurrments, with the chlorides, mineral were determined before and after the exgave 1.362 -I: 0.009 meq./g. As would be supposed, the posure t o the acidified solutions. The values obnature of the anion is without effect. tained (1.008 meq./g. and 0.980 meq./g., respec-
t
cl
L
-
Resulls The exchange isotherms are given in Fig. 1 and 2 . These are deviation Plots; the fraction of sites occupied by cesium ('l/din excess of the equivalent fraction of cesium in the solution is given as the ordinate and the equivalent fractiol? .of cesium in solution (c/co) as the abscissa. Positive values thus indicate selectivity for cesium. The complex nature of the results is quite evident. I t might be remarked that the work with cerium was done later and entirely independently of the work with yt,trium. The qualitative agreement between the two cases supplies a desirable check. The results given in Fig. 1 and 2 were obtained with solutions initially adjusted t o about pH 6.8 and without further attempt at pH control. To determine the possible effect of small variations in pH, a series of experimeuts in the cesium-yttrium system was carried out in which the pH was deliberately varied over a range of nearly four units (as measured with a glass electrode) by the addition of small amounts of concentrated nitric acid to the solutions. The pH of the effluent equilibrium solution differed from that of the input solution by no more than 0.15 pI-1 unit. These results are given in Table I. The experiments were done in duplicate or triplicate, and from these data the character of the precision obtainable in this work
tively) indicate the possibility of a small loss due to acid attack. No trends, however, are observable in the extent of adsorption. To withill the precision of these experiments the exchange sorption of yttrium-cesium is entirely unaffected by the pH of the solution in the range 3-7. It is clear that whatever may be the variation of the yttrium species it1 the solution the proportion of yttrium on the sorbent is unaffected by changes in p H . We assume that cerium exhibits similar behavior. TABLE I ~ p p E c TOF
p~
O N THE C E S I U ~ I - ~ T T R I U MEXCHANGE ON
MONTXORILLONITE Temp. 50"; ca = 0.04 N c/co
(CS)
0 30 0 TO
pEI (30')
6 4 3 6 4 3
8 5 2 0 5 2
d P 0 (CS)
0 367,O ,378, 374, 814, 814, 811,
378,O 364 386 380 822,O 824 814 805
The agreement with the data of Fig. 211 is eiitirely satisfactory at c 'eo = 0.70; at c co = 0.30, on the steep part of the isotherm, the agreemelit is perhaps as good as could be expected. At low total concentration the clay is selective for the trivalent ion both when this ion is present
GALENR. FRYSINGER AND HESRYC. THOMAS
226
as to the charge number a series of ten columns was run, Y* + (natural clay), Yo --t Y*, a t 75". The measured capacity (assuming tripositive Y, 1.108 & 0.006 meq./g. clay as weighed) corresponds to a charge number of 2.8. This result must be considered more accurate than that from the double tracer experiments. The self-consistency of these results shows clearly that whatever may be the nature of the species in the nearly neutral yttrium solutions, the clay prefers its trivalent ions a t the
a
q;
+0.200
0.100
0.000 +0.200
-; /
/dL>
j4
Vol. 64
-
~~
I
~~
I
1
1
1
[
18'
1
,-. u"
\
2 +0.100 I A
d
3
0.000
+0.100
I
/
0.000
-0.100
\\ \
!
I
l
,l
0.4
0.2 ,
I
,,
0.6
\
I
0.8
by
I
\?:/
\
\ c/ eo).
Fig. 2.-Exchange adsorption isotherms, Cs-RE, on montmorillonite a t 0.t4 N : 0,yttrium; 6 or 9 cerium, (a) 30"; (b) 50'; ( c ) 75 .
in small proportion and when it is present in large proportion. In the middle range of composition the clay selectively sorbs cesium, except at high temperatures and low concentrations; this selectivity decreases rapidly with the temperature. Certainly in the case of yttrium at low total concentration the maximum selectivity for cesium shifts toward higher cesium content with increase in temperatures. The data for cerium are not sufficiently detailed to show if this effect also occurs with this element. All of the data presented in Fig. 1 and 2 are based on determinations of cesium via the displacement Cso + Cs*. The reproducibility of these results as determined from independent experiments designed to be identical is about 0.006 in q/qo regardless of its magnitude. In one series of experiments with yttrium both metals were traced. Knowing the hold-up volume of the column and assuming a constant capacity of the clay, we can calculate from the results with these "double-isotopic" columns the average charge number of the sorbed yttrium. Errors accumulate badly in this computation, but the results point clearly to an increase in the extent of hydrolysis of the sorbed trivalent metal with increasing temperature. A series of six experiments (at different cesium contents) gave the following charge numbers: 30", 2.96 f 0.16; 50", 2.84 =k 0.15; 75", 2.65 f 0.14. (The maximum differences from the average are given.) There is no clear evidence of an effect due to cesium content. To check this conclusion
in which M and N refer to molalities and fractionof-sites for the indicated substances; in the second form all quantities refer to cesium, and molarities and molalities are identified for the dilute solutions in question. The calculation of the thermodynamic equilibrium constant is summarized in the formula
in which
and thef's are the activity coefficients in the monoion clay phases at the molalities in use. Thef's are defined as unity in the mono-ion clays in equilibrium with infinitely dilute solutions. These situations are takcn as the standard states of the solid phases. We dispose of the complications in formula (2) as follows: According to (3) we need measurements of the mean activity coefficients in mixed solutions of the rare earths and cesium. These are not available. In pure cesium nitrate solutions measured activity coefficients are available8 only down to 0.1 M , and we can find no data for pure yttrium nitrate. Therefore we use the Debye-Huckel expression for the activity coefficient, taking the a parameter fpr CsN03 equal to that for CsC1, 2amely 3.0 A.9; for both rare earths we use 4.95 A. as given by Glueckauflo for yttrium. The Gug(7) G.L. Gaines, Jr., a n d H. C. Thomas, J . Chcm. Phys., 21, 714 (1953). (8) R. A . Robinson and R . H. Stokes, "Electrolytic Solutions," Butterworth, London, 1955, p. 480. (9) H. S. Harned a n d B. B. O n e n , "The Physical Chemistry of Electrolytic Solutions," Reinhold Publ. Corp., N e w York, N. Y . 2nd Ed., 1950, p . 381. (10) E. Glueckauf, Nature, 163, 4lk (1949).
Feb., 1960
YTTRIUM-CESIUM AND
CERIUM(III)-CESIUM O S h~ONrMORILLOSITE
genheim equation as modified by Glueckauf gives essentially identical results at 25”, and since the data for using this equation at other temperatures are lacking, no further attempt was made to improve upon the simple Debye-Huckel calculation. The computations of the activity coefficient ratios mere made for the various compositions at the three temperatures. The contribution of this term to (2) is far from negligible; it amounts to about 1.1in In
227
-2
-1
K. The last term in (2) gives the contribution to the free energy due to the change in water activity in going from (a) (that is, CON, Y(NO3)3) to @)(CON, CsN03). Since detailed data on the water compositions of t)heseclays are not available, we write
where ~ Z H ~isOan average water content in moles per -1 exchange equivalent. A number of investigat o r ~ ” - ’ have ~ made studies of the montmorillonite-water system, from which we may get estimates for f i H z 0 . With increasing water activity the c-axis spacing of montmorillonite increases by integral steps, corresponding to mono-layer incre- -2 ments of water molecules. This “crvstalline swelling” brings the spacing to about 20 A. for 99% humidity, corresponding to four layers and a total 0.2 0.4 0.6 0.8 interlayer water content of about 0.5 g./g. clay. O/Oo ((3s). _ _ No direct experimental data are available for the Fig. 3.--Solid phase activity coefficients, 30”, 0.02 AV. water content of the montmorillonites in our solutions. Norrish,’s however, gives a c-axis spacing of The resulting expression (approximate) for In 13.8 A., for a cesium montmorillonite in equilibrium [ f ~ ~ ( a ) / f c ~cancels ~ ( b ) ]the contribution of (4) in In with water and 19.4 8. for aluminum montmorillo- K . These “zero-point” contributions must, hownite in equilibrium with 0.01 N A12(S04)3. These ever, be retained in the computation of the solid spacings correspond to a water content of 0.3-0.4 g./ phase activity coefficients. When the small term g. clay. I n the absence of any better data we take due to the variation of the water activity is ne0.5 g. as an upper limit for the average water con- glected, these activity coefficients are given by, e.g. tent. This corresponds to f i ~ ~ o26.6 moles per exchange equivalent. We calculate the ratio of the activities of the solvent in the two electrolytes from the osmotic coefficient, using the Debye-Huckel ap(6) proximation. The contribution of the water term is, of course, small for our dilute solutions, the larg- Complete calculations of the equilibrium constants est value being 0.049 in In K (at 0.04 N , 75”). I n and activity coefficients have been made for ytany case, in the approximation to which we are trium and for cerium.l8 The general behavior of forced to work. the contribution will be seen to can- the activity coefficients is illustrated in Fig. 3, in which as is seen, the complicated nature of the adcel. The first non-constant term in ( 2 ) gives the con- sorption process is reflected in the relatively enortribution t o the standard free energy of the process mous deviations from unity of these quantities. In Table I1 the values of the standard free enerof bringing the mono-ion clays from their standard states into equilibrium with solutions of normality gies are given for the different temperatures. cfl. I n the model we use, the effects, if any, of TABLEI1 sorbed anion are entirely disregarded. This activSTAKDARD FREEENERGY CHANGES FOR THE REACTIOXS ity coefficient ratio is then determined entirely by the properties of the solvent. Neglecting the moCs&l ( R E ) + + + ( N )= 3Cs+(N) (RE)M AF@(kcal./moleRE-j lal volume of the solvent in comparison with that of ---Yttrium----Ce!iun----. the vapor and assuming, perforce, a constant water T 0.02 N 0.04 Ai 0.02 S 0.04 N content, rye find 303 4.81 5.01 4.49 4.68
