Anion Exchange of Metal Complexes.
X1V.l
The Effect of
Acidity on the Sorption of Lanthanides from Lithium Nitrate Solutionsz
by Y. Marcus and M. Givon Radiochemistry Department, Soreq Research Establishment, Israel Atomic Energy Commission. Yavne, Israel (Recehed March 6 , 1964)
Activity coefficients of sniall concentrations of nitric acid in 1-8 nz lithium nitrate have been deterniiiied poteiitionietrically, using quinhydrone and glass electrodes. The invasion of a Dowex 1 resin by nitric acid and lithium nitrate from a 3.90 m total nitrate solution has been measured. Activity coefficients of nitric acid in the resin were calculated. The acidity dependence of the distribution of La, Eu, Yb, and tAm between resin and a lithium nitrate solution was measured. Sorption of the species ;\I(SOs)6-2,for RI = Eu, Yb, and possibly Ani, and La(N03)1-4was found to be consistent with the experimental data.
Introduction It has been noted3 that the concentration of nitric acid has an appreciable effect on the sorption of a number of lanthanides on an anion exchanger from lithium nitrate solutions. Such an effect has also been observed for the sorption of americium’ and other ions.4 It is generally observed that an increase of the concentration of nitric acid a t constant total nitrate concentration causes a decrease of the distribution coefficient. I n order to give a quantitative interpretation of this effect, additional information is required. This includes the activity of the nitric acid in the lithium nitrate solution and in the resin. These data were obtained from series of experiments, in which the potentials of quinhydrone and glass electrodes were measured as a function of the acidity a t constant nitrate concentration, with and without resin present. Experimental Materials. Reagent grade chemicals were used throughout. The resin was Dowex 1-X 8 for the acid and americium sorption experiments and Dowex 1-X 10 for the experinients involving sorption of the lanthanides. Quinhydrone was once recrystallized from water. The americium was a sample of nominally Am241,obtained from Oak Ridge Kational Laboratories. Methods. Potentiometric measurements were carried out a t 30” using the cells T h e Journal of Physical Chemistry
Pt,’QHZ, Q, LiKOs(mL,), HNOs(mH)/Is.c.e. (1) GE,’LiXOa(mL,), H y O g ( m ~//s.c.e. )
(2)
where QH,, Q denotes saturated quinhydrone, // s.c.e. the saturated calomel electrode with a liquid junction, and GE a glass electrode. The instrument used was a Radiometer 22 pH meter, provided with a GK 2021 combined glass-calomel electrode. Solutions were stirred magnetically. An attempt to use a platinumblack-hydrogen electrode resulted in unsteady potentials in nitrate solutions, contrary to its behavior in chloride solution^.^ Acid was added to a lithium nitrate solution from a microburet. Resin in nitrate form was prepared by treating the chloride form in a coluinn with a large excess of lithium nitrate solution. The water content of the wet resin was determined by the centrifugation method, that of air-dry resin by heating in a vacuum oven at 60”. Imbibed nitrate was determined by elution with water and spectrophotometric measurement at 300 mM ( E 7.15). The capacity (2.95 mequiv./g. of dry resin) was deter(1) Previous paper in series: Y . Marcus, h l . Givon, and G. R. Choppin, J . Inorg. .Vucl. Chem., 2 5 , 1457 (1963). (2) Presented in part at the 1st Meeting of the Israel Chemical Society, Beer Sheva, December, 1963; Israel J . Chem., 1,255 (1963). (3) Y . Marcus and F. Nelson, J . Phys. Chem., 63, 77 (1959). (4) F. Selson and K. A . Kraus, J. Am. Chem. Soc., 76,5916 (1954); R. F. Buchanan and J. P. Faris, Conf. Use Radioisotopes Phys. Sci. Ind., Copenhagen, 1960, paper RICC/173. (5) R. G. Bates, “Electrometric pH Determination,” John Wiley and Sons, S e w York, N. Y . , 1954.
2231
ANIONEXCHANGE O F METALCOMPLEXES
mined for the chloride form by elution with nitrate and potentiometric titration with AgNOa. The distribution of the lanthanides was measured at 25 and 78" by the column technique, as reported previously3; that of americium was measured at ca. 25" by the batch method, analyzing the solutions both spectrophotometrically and radiometrically.
Table I : Some Values of ~ H Q~ ,H Q Eand ,
m"oa
x
mLi =
1.00
Results and Discussion A . Nitric Acid Actzuities in Lithium Nitrate Solutions. The e.m.f. values of cell 1 were measured for mL1 1.00 m, 1.99 m, 3.98 m, 5.55 m, and 8.00 m for various acid concentrations, ranging from about m to lo-' m. (The acid concentration is very small as compared with the lithium nitrate concentration, which is therefore practically equivalent with the total nitrate concentration.) The e.m.f. values were corrected for the liquid junction potential and salt effect on the quinhydrone electrode according to the equation
YH
a t Various LiNOa Molalities
=i
mLi = 3 . 9 8
108
PHQ
1.12 2.23 4.46 8.35 16.7 33.4 55.7 112
2.73 2.50 2.23 1.95 1.68 1.38 1.18
1.22 2.42 4.84 9,67
2.34 2.14
18.2 36.2 60.4 122
0.90
1.81 1.51 1.21 0.90 0.68 0.40
APH
7H
1.66 1.41 1.33 1.34 1.24 1.24 1.19 0.00 -0.01 1.13 YH = 1.33" 0.02 0.02 0.02 0.05 0.03 0,02
0.16 0.19 0.19 0.19 0.24 0.22 0.24 0.25 YE =
The liquid junction correction AEj between lithium nitrate and saturated potassium chloride was calculated from the Henderson" equation, using equivalent conductances from standard compilation^.^ The salt effect correction AE,,lt was calculated according to StonehilL8 The value for lithium nitrate was obtained from a linear combination of the values for potassium nitrate arid chloride and lithium chloride. The correction was found to be AEi,&lb = 3 . 8 C ~ imv., and the appropriate molal values were calculated from the equationg
The total corrections were 7, 14, 25, 33, and 45 mv., respectively, in increasing order of the lithium nitrate concentrations used. From the e.m.f. data (in mv.) pH values ( ~ H Qwere ) calculated as a function of mH
where (Eao -- E,,,') = 454.0 mv. at 30" was taken from Bates.6 The results are shown in Table 1. Mean molal activity coefficients of nitric acid in lithium nitrate solutions (rrt" 0 % (In L ~ N o ~ ) desig, nated in short YH in the following) were calculated from the PHQ data a s a function of the total nitrate concentration. They, too, are shown in Table 1. It can be seen that YH does not vary appreciably with the acid concentration. The average values of Y H are plotted in Fig. 1 as a function of the lithium nitrate concentration, together with values of the mean ac-
rnLi =
8.00
1.32 2.64 5,28 9.90 19.9 39.6 66.0 132
1.53 1.15
0.87 0.60 0.29
0.00 -0.18 -0.44
0.31 0.35 0.36
0.38 0.40 0.41
0.40 .. YE
a
3.68 2.99 3.20 3.20 3.40 3.47 3.47 3.29 3.31"
15.4 18.4 17.7 17.5 17.8 17.3 15.8 14.4 = 17.7"
Mean value of Y B (including data not shown in the table).
tivity coefficient of lithium nitratelo and of the ionic part of nitric acid.ll The activity coefficients of nitric acid in lithium nitrate are seen to be much higher than both Y + I , ~ N O ~ in lithium nitrate and ~ Y H + Y N O ~ - . A similar high result which is also shown in Fig. 1 was obtained by Rosenthal and Dwyer12 for nitric acid at low concentration in 6 M (about 7.5 m) sodiuiii nitrate. The data of Fig. 1 lead to values of o!H for Harned's'3 rule
(6) P . Henderson, 2.physik. Chem., 59, 118 (1907). (7) B. E . Conway, "Electrochemical Data," Elsevier Publishing Co., Amsterdam, 1952; R. Parsons, "Handbook of Electrochemical Constants," Butterworth and Co., Ltd., London, 1959. (8) H. I. Stonehill, Trans. Faraday SOC.,39, 67 (1943). (9) Y. Marcus and I. Abrahamer, J . Inorg. Nucl. Chem., 22, 141 (1961). (10) R. A. Robinson and R. H. Stokes, "Electrolyte Solutions," Butterworth and Co., Ltd., London, 1959. (11) H. A. C. McKay, Trans. Faraday SOC.,52, 1568 (1956). (12) D. Rosenthal and J. S.Dwyer, J . Phys. Chem., 66,2687 (1962). (13) H. S. Harned, J . Am. Chem. Soc., 48, 326 (1926).
Volume 68, Number 8
August, 1964
2232
Y. MARCUSAND M. GIVON
log
YH
=
log Y*Hivo,(mLi) -
C Y H ~ L ~
(6)
which range from -0.26 at 1 m to -0.12 at 8 m lithium nitrate, Le., very high values. If Guggenhejm’s treatment14 were applicable, LYH a L i should be zero, and
+
I
-
I
I
0.51
I
I
I
0
2
4
6
I
I
I
a NO;
Figure 1. Mean molal activity coefficients in nitrate solutions: 0, Y + L ~ N O in ~ lithium nitratelo; 0 , ~ \ / Y E + Y N O ~ for nitric acid in nitric acid”; A, Y I E N O ~ in lithium nitrate, ; Y ~ H N Oin~lithium nitrate from this work, from ~ H QA, PHGEand eq. 7; 0,Y A H N O ~in sodium nitrate.’$
- a L i = ~ ( ( P H N O ~- ~ ~ i ~ o , ) / 2 , 3 0 3where m ~ i the v-values are the osmotic coefficients for the pure components of total molality m ~ i , Values of a~ calculated by this method, using ~ L ~ N Ofrom , ref. 10 and ( P H N O ~ calculated from the water activities, yielded values about eight times smaller than the observed values. This means that the hydrogen ions are influenced not only by the nitrate but also by the lithium ions.14 Parallel determinations of the pH values were made for 1.0, 2.0, 3.0, 4.0, 5 . 5 , 6.5, and 8.0 m lithium nitrate solutions with a glass electrode calibrated against two standard buffers of nominal pH 4 and 7. The data ( ~ H G Eare ) given in Table I. It can be seen that for each lithium nitrate concentration the differences Q only slightly with the acid ApH = PHCE - ~ H vary concentration and are a linear function of the total nitrate concentration. (YH
ApH
=
The Journal of Physical Chemistry
0.048m~i
(7)
This difference in pH is probably due to the lithium ion effect on the glass electrode. The functions pHGE us. acidity were used as empirical means for determining the actual acid concentration from glass electrode measurements. B . Nitrzc Acid Activity in the Resin Phase. Two series of experiments were made in which small known amounts of 4.04 m nitric acid were added to a known quantity of 3.90 m lithium nitrate solution containing a weighed amount of anion-exchange resin, and the p H was measured with a glass electrode after attainment of equilibrium. The acidity range in solution was 3 x to lo-’ rn. The results were corrected for dilution and it was assumed that the pH values which in this instance were measured a t 25” could be used with the empirical curves which were determined at 30°1 From these measurements the equilibrium acidities in solution were calculated and from the difference between initial and equilibrium acidity the invasion into the resin was estimated as molality of nitric acid in the resin %H. For this purpose values of the equilibrium water content of the resin are required. These were determined by the centrifugation method for resin in equilibrium with a practically neutral lithium nitrate solution and for the resin a t the end of each run. Water contents of (410 mg./g. of dry resin) did not vary appreciably with acidity. The invaded electrolyte was eluted from the resin a t the end of one run and the acid and total nitrate concentration determined directly. Values of f i us. ~ mH are shown in Fig. 2. The point for the directly determined acidity falls neatly on the plot of r i i ~us. mH which is linear with a slope of 19.5 f 1.5. The total concentration of the invading nitrate in the resin was 4.96 m, of which 1.72 m was nitric acid and 3.24 m lithium nitrate, The latter value agrees well with the value interpolated from Danon’s16data, which shows that it is independent of the acidity, The niolality of exchange sites is the ratio of the capacity to the water content, ”i25 m, and the total nitrate molality in the resin is 12.2 m. From these molalities the ratios of the mean activity ~ r H = coefficients in the two phases, r L 1 = ~ L , / Y L and T H / Y H (where TM and Y M designate in short Y ~ M M N O ~ in the resin and solution phases, respectively) are calculated from the relations
rL12 = mLlmxo,/fiL1(mL1+ “ j L 2 ~ + m
~ (sa) )
and
rH2= mHmNo,/mH(mL1 + f
+
i ~ f i ~ (8b) )
(14) E. A. Guggenheim, Phil. Mag., 19, 688 (1935). (15) E. Hogfeldt and B. Bolander, Arkiv Kemi., 21, 161 (1963). (16) J. Danon, J . Phys. Chem., 6 5 , 2039 (1961).
ANIONEXCHANGE OF METAL
1.5
-
1.0
-
2233
COMPLEXES
approximately 0.13, as conipared with 0.36 for resin in equilibrium with pure 3.9 m nitric acid.16 It is seen that high concentrations of lithium nitrate decrease rH considerably. Since TH is a measure of the relative acid strength in the resin and the solution, it is clear that the invading excess lithium nitrate causes a decrease in the acid strength. This result may be compared with the data of Xelson and Kraus” for hydrochloric acid in lithium chloride at a total molality of 9.9. Hydrochloric acid was found to be rather weak in the presence of high concentrations of lithium chloride, TH = 0.14, compared with TH = 0.60 for pure 9.9 m hydrochloric acid. Again TL, changes only little; it increases (contrary to the nitrate case where it decreases) from 0.77 to 1.09 as the acid concentration increases. The decrease in acid strength in the resin may be explained by the hypothesis of the formation of the binitrate ion H(r\T03)2-. Such an ion could be formed by ion-dipole interaction in low acid-high nitrate solutions. The low effective dielectric constant of the resin phase could favor its formation. I n a medium of still lower dielectric constant, such as toluene18 or a solution of a trialkylamine in a hydrocarbon, l9 the dimerization of nitric acid has been found. For the corresponding chloride system a species HC12- has been proposed.20-22 Formation of this H(XO3)z- species in the resin will cause a decrease in the concentration of free nitrate ions as the nitric acid concentration in the resin increases. If all the acid is in this form, then
0.5E0
0.0 0.00
0.02
0.06
OD4
0.08
HNO,
0
Figure 2. Xitric acid invasion into Dowex 1-X 8 from 3.90 m lithium nitrate as a function of the acidity: 0, from ~ H G E0;, direct determination.
The results are shown in Table 11. It can be seen that T L decreases ~ from 0.67 to 0.62 as the acidity of the solution increases from 3 X to 9 X m. The value a t low acidity is consistent with the value of 0.68 interpolated from Danon’sl8 data. The value of r H is
Table 11: Values of
p H and FL, for Dowex I - x 8 a t 3.90 m Nitrate Ion Concentration, %L, = 3.24, %R = 7.22
mH
0.00028 0.0047 0.0063 0.0089 0,0123 0.0180 0.0250 0.0315 0.0385 0.0405 0.0510 0.0520 0,0628 0.0630 0.0740 0.0840 0.0910
(GH
-
GLi
mH
mLi
+ mR)
rH
rLi
0.014 0.15’2 0.250 0.340 0.330 0.470 0.540 0.630 0.670 0.670 0.870 0.940 1.270 1.121) 1,400 1.170 1.290
3.90 3.89 3.89 3.88 3.88 3.87 3.86 3.86 3.85 3.85 3.84 3.84 3.84 3.84 3.83 3.82 3.81
10.47 10.61 10.71 10.80 10.79 10.93 11 .oo 11.09 11.13 11.13 11.33 11.40 11.73 11 58 11.86 11.63 11.75
0.086 0.106 0.096 0,097 0.116 0.117 0.129 0.132 0.141 0.144 0.140 0.135 0.125 0.135 0.129 0.151 0.150
0.670 0.665 0.660 0.655 0.658 0.655 0.655 0,650 0.650 0,645 0.640 0.635 0.635 0.630 0.630 0.625 0,625
?%NOa = ?%R
?%Li
-
?%H
(9)
C . Distribution of Lanthanzdes and Americium as a Function of Acid Concentration. Results for lanthanum, europium, and ytterbium obtained previously2 in 4.6 m lithium nitrate and values for americium in 4.0 m lithium nitrate are shown in Fig. 3. The curves are normalized against the distribution coefficient a t very low acidity, Do. Some experiments for lanthanum and europium were carried out a t both 25 and 78”, but no significant differences between the results were found. The figure shows that above about 0.01 m nitric acid the distribution coefficients begin to be affected by the acidity. (17) F. Nelson and K. A. Kraus, J . Am. Chem. Soc., 80,4154 (1958). (18) C. J. Hardy, B. F. Greenfield, and D. Scargill, J . Chem. Soc., 60 (1961). (19) E. Hogfeldt and F. Fredlund, private communication. (20) S. Lindenbaum and G. E. Boyd, U.S.A.E.C. Report ORNL3320, 1982, p. 75. (21) J. R. Beatty and G. J. Leigh, J . Chem. Soc., 4726 (1962). (22) G. Duyckaerts, J. Fuger, and W. Muller, Euratom Report No. 426.f, 1963.
Volume 68, Number 8 August, 1964
DO log -
D
=
p log NO^p log
LH
?mLl
-4.0
-3.0 I
I
I
I
-1.0 logmHNO,
-2.0
Figure 3. Acidity effect on the distribution of ( 0 )La, (m) Eu, (A)Yb, and ( I ) Am between lithium nitrate and anion-exchange resin. The curves are calculated from eq. 12 with p = 2 (lower) and p = 4 (upper).
The following argument can be used to interpret the data. It is assumed that the metal ions distribute between the resin and the solution according to the reaction
It has been shown23that the distribution coefficient is proportional to the pth power of the effective nitrate activity in the resin. This latter quantity is the product of the free nitrate ion concentration and the effective activity coefficient of the nitrate ion28 which is i TH accordassumed to vary linearly between y ~ and ing to the fractions of salt and acid invading the resin
Since in the absence of acid the effective ligand activity would be (@m f i ~ ) y L i ,the ratio of the distribution coefficients in absence and in presence of acid will be
+
The Journal of Physical Chemistry
+
p log (PTZR
rTZLi)jLi
iree)TN08eff
(GZR
+
a L i
-
-
= P log %H)
&H
+
TLi
*LL
+
*LJTLi
-
X
*Li +
(*R
1
=
~ F H (12)
where F H is a correction function. Values of F H for 3.9 m lithium nitrate were calculated from the data obtained in parts A and B above. The various random errors in the quantities entering the calculations and systematic errors such as the difference in temperature between the data in parts A and B and in concentration between the data in parts B and C were estimated to yield a relative error of *20% in F H a t the 95% confidence limit. Within this error limits, eq. 12 is seen to fit the results for the lanthanides (Fig. 3) with p = 2 for europium and ytterbium and p = 4 for lanthanum. Application of eq. 12 with p = 3 does not fit the data. The data for americium are not sufficiently precise for a good fit. If it is accepted that the fit in Fig. 3 is evidence for the operation of the acid effect according to eq. 12, and absorption on the resin according to eq. 10, then the values p = 2 and p = 4 indicate the formation of E~(,ly’o~),~-, Yb(N03)62-,and La(S08)74-. High coordination numbers for the lanthanides, up to 8 or 9, have been variously reported,24so that the high values of the ligand number do not seem to be improbable, provided the nitrate is a monodentate ligand. Species with five nitrate ions per lanthanide have been shown to form in long-chain amine extractionsg A species Nd(N03)6’- has been found in a spectrophotometric study of Kd(N03)ain ethanol.*,
Acknowledgment. Thanks are due to Nessrs. A. Gafni and M, Compan and Mrs. E.Bauman for help in carrying out the experiments. (23) Y. Marcus and C. D. Coryell, Bull. Res. Council Israel, 8A, l(1959). (24) E.g., Nd(HC0s)en- and Eu(HCOs)74-, J. A. Marinsky and H. S. Sherry, Inorg. Chem., 3, 334 (1964); cf. also G. Vincentini, J . Inorg. Nucl. Chem., 24, 1351 (1962); P. W. Selwood, “Magnetochemistry,’’ Interscience Publishers, Inc., New York, N. Y., 1956, pp. 152-153. (25) I. Abrahamer and Y . Marcus, Israel AEC Semi-annual Report, June-December, 1963, IA-920 (in press).
