of the solution by mounting an 10O-bl. aliquot for gross a- and gross @-counting. Determine the lithium concentration of the solution by flame photometry. (The working curve should be prepared by analyzing standard solutions of lithium in 0.1V ammonium hydroxide0.005N ammonium nitrate.) The calcium fluoride column may be used for a t least three more sample aliquots. RESULTS AND DISCUSSION
The mechanism for the removal of trivalent lanthanides and actinides from solution by calcium fluoride and the factors affecting the performance of the calcium fluoride column have been described ( I ) , The highest decontamination factors are obtained when the sample and washed are basic. The wash solution, 0.1N ammonium hydroxide, is made 0.005M in ammonium nitrate, which serves as an electrolyte to prevent peptization of calcium fluoride and breakthrough of radioactivity. Four synthetic samples, each containing 1.2 x 1 0 8 d.p.m. of americium-
241 alpha activity, were passed consecutively through a single calcium fluoride column. Because the respective decontamination factors, (a d.p.m. in feed)/(a d.p.m. in effluent), were 7.5 X lo6, 5.0 X lo6, 2.5 X lo6, and 7.5 X lo6, a column can be used for at least four samples. The degree of separation of lithium from the trivalent actinides and lanthanides was determined with 16 simulated Tramex feed solutions, each containing 1.2 x 108 a d.p.m. of americium-241. A separate calcium fluoride column was used for each sample. The average decontamination factor was 5 X 106. The effluent and wash solutions from seven of these samples contained less than detectable activity representing complete removal of activity by the calcium fluoride column. Twelve simulated Tramex feed solutionq were passed through four calcium fluoride columns (three samples per column) and were analyzed for lithium by flame photometry. Each of the synthetic samples was 11. O M in lithium.
The results of the twelve analyses averaged 10.7-11 lithium (97.3% recovery) with a relative standard deviation of 1.7% ( n = 12). One actual feed s d u tion containing 10' CY d.1i.m. and 1.5 X lo7 p d.1i.m. per sample aliquot was analyzed for lithium with a relative standard deviation of 1.8% ( n = 5). This method is being used routinely a t the Savannah River Laboratory. ACKNOWLEDGMENT
The author acknowledges the assistance of E. G. Orebaugh, who performed the flame photometric analyses. LITERATURE CITED
Holcomb, H. P.,
A . v . 4 ~ . CHEM.36, 2329, (1964). ( 2 ) Leuze, R. E . , Bavbarz, R. I)., Weaver, Boyd, Yucl. Sci. Eng. 17, 252 (1963).
(1)
H. PERRY HOLCOMB Savannah River Laboratory E. I. du Pont de Semours & Co. Aiken, S. C. WORK done under contract AT(07-2)-1 with the U. S. Atomic Energy Commission.
Kinetic Inhibition of Cerium(lV) Reduction by Oxalic Acid SIR: I n a previous paper ( 2 ) we reported on a detailed kinetic study of the cerium(1V)-oxalic acid reaction in sulfate media t o provide basic information useful in choosing optimum reaction conditions for analytical titrations. Our study of this reaction has now been extended t o conditions which prevail a t the beginning of a titration-the situation where one of the reactants is present in great excess. The primary aim of the experiments reported here is to further confirm the proposed mechanism ( 2 ) and, if possible, to explain the anomalous induction period described by Rao, Mohan, and Sastri (3) for the first addition of cerium(1Vj to excess oxalate in sulfuric acid media. Using the classification scheme of Anbar ( I ) , 01 r data for the mechanism fit the sequence MX,+a (stable) Lj 2X J I X , - l L + ~ (unstable]+ MX,-l +h products where Ai', X, and L stand for cerium(IV), sulfate, and oxalate, respectively. According to Taube (4, 6), a likely consequence of this mechanism, if n > 1, is that the specific rate of the oxidationreduction reaction should increase, pass through a maximum, and again decrease as the ratio of I, t o JZX,+a is increased. We tested this hypothesis,
+
+
+
using the techniques described in the first paper ( 2 ) )by following the kinetics of cerium(1V) disappearance in solutions containing a 2- 120-fold excess of oxalate under otherwise identical conditions. The results are summarized in Table I and show that the specific rate constant for the cerium(IV)-oxalate react:on indeed passes t,hrough a maximum and then decreases as the ratio of oxalate to cerium(1V) is further increased.
Table I. Effect of Excess Oxalic Acid on Specific Rate Constant 4 08 x 10-4M Ce(IV), 0 5M H280r, 5.0 + 0.05" C.
Oxalic acid, M 0 0 0 0 0 0 0
050 025 020 010 005 002 001
Specific rate constant, sec. -l 0 0 0 0 0 0 0
022 043 093 093 085 049 021
The most probable mechanistic explanation for the observed trend involves the formation of higher
oxalato complexes of cerium(1V) by successive replacement of sulfate ligands by the reducing agent. In selective oxidations of this type such complexes tend to be less reactive than the 1 : l intermediate (5). The oxalate to ceriuin(1V) ratio a t maximum specific rate reflects the formation constant of t'he intermediate and is consistent with independent optical measurements ( 2 ) . The principal analytical consequence of this kinetic complication will be an apparent induction period a t the beginning of titrations where crrium(1V) is titrated into excess oxalate. R'e believe that the anomalous results described qualitatively by Rao, Mohan, and Sastri (3) are caused by t,his effect. LITERATURE CITED
( 1 ) Anbnr, XI., Surnnier Symposium on
Inorganic lfechanisms (ACS), Lawrence, Kan., June 1964. ( 2 ) El-Tantawy, Y. A., Ilechnitz, G.A , , ANAL.CHEM.36, 1774 (1964). ( 3 ) Rao, (>. P., lfohari, 1'. G.,Sastri,
&I. S . , 2. =Lnal. C h e m . 156, 338 (1957). ( 4 ) Tauhe, H . , .J. L A t u . Chem. SOC. 69, 1418 (1947). (5) Ibid., 70, 1216 (1948). G. A. RECHNITZ Y. A. EL-T.4NTAR.Y
Universit . of Pennsylvania Philadelpha, Pa. VOL. 36, NO. 12, NOVEMBER 1964
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