Jan., 1957 THERMAL ISOTOPIC EXCHANGE BETWEEN ETHYL IODIDE AND MOLECULAR IODINE IN THE LIQUID PHASE1 BY L. R. DARBEE A N D G. M. HARRIS Contribution /Tom the Department a/ Chemistry of the University of Buffalo,Buffalo,N . Y . Received J u l y 10, 1966
Several investigators have reported esperiments on the isotopic exchange reaction between iodine and ethyl iodide in the liquid p h a ~ e . ~ -On ~ the whole, the results have been inconclusive insofar as determination of the kinetics of the process are concerned. The result of a further attempt to elucidate this mechanism is reported herewith. Experimental Purifications of Ethyl Iodide.-The best obtainable grades of a number of commercial ethyl iodides were treated to remove traces of impurities such as water, alcohol, iodine and hydrogen iodide. The “best” procedure, as judged from the results of preliminary experiments to be described below, was as follows: After successive washings with aqueous sodium thiosulfate and distilled water, the iodide was dried over anhydrous sodium carbonate, filtered through glass wool, and fractionated just prior to use in a 23 theoretical plate ceramic-packed column. A final trap-to-ttrap distillation preceded storage under vacuum. Preparation of Radioactive Iodine.--Inact#ive solid iodine, purified as suggested by Pierce and Haenisch,R was added to aqueous carrier-free N a P . The mixture was acidified inactive NaI added.’ with H1S04,and a small amount of C.P. Iodide was oxidized to iodine with KMnOd, the aqueous phase withdrawn and discarded, and the residual iodine containing nearly all the activity washed with water and dried by distillation through a glass wool-Pz06 mixture to storage under vacuum.8 Technique of Exchange Runs .-Active iodine was metered by vacuum technique into sample tubes containing measured quantities of carefully outgassed ethyl iodide cooled in liquid nitrogen. Series of identical sample tubes (lengths of 9 mm. Pyrex tubing) were sealed off under high vacuum, wrapped in aluminum foil to exclude light, and suspended in a thermostat maintained at the desired temperature to within &O.O5O, from which they were withdrawn singly a t intervals for assay. Fixed aliquots of each sample were diluted to a standard volume with heptaneg and placed in a 25-ml. erlenmeyer flask equipped with a metal collar enabling it to be positioned reproducibly over an inverted, cylindrically-shielded, end-window G/M tube. This count gave total activity. The heptane solution was then extracted with aqueous thiosulfate, enabling counts to be made of “inorganic” and “organic” activities. Low iodine concentrations in pure ethyl iodide solution were determined by its absorption in the 4700-4900 A. region. Ethyl iodide a t higher iodine concentrations required dilution of the mixture with known proportion of heptane before spectrophotometry. Beclcman D U and DK2 spectrophotometers were employed. I n experiments where ethanol was added, re( 1 ) From the doctoral dissertation of L. R. Darbee, University of Buffalo, 1956. Complete text available from University Microfilms, Ann Arbor, Michigan. (2) D. Hull, C. Shiflett and S. C. Lind, J . Am. Chem. Sac., 68, 535 (1936). (3) P.F. D.Shaw and C. H. Collie, J . Chem. Soc., 1217 (1949). (4) R. Noyes, J . Am. Cham. Soc., TO, 2614 (1948). (5) R. G. Badger, C. F. Chmiel and R. H. Schuler, ibid., T S , 2498 (1953); C. F. Chmiel, Master’s Thesis, Caniaius College, 1952. (6) W. Pierce and E. Haenisch, “Quantitative Analysis,” 2nd Ed., John Wiley and Sons, Inc., New York, N. Y.,1940, p. 203. (7) This promoted reduotions of 11*108-, aometimes present in aged carrier-free solutions. ( 8 ) Of several stopcook greases tried, “Ascolube” showed the greetest resistance t o attaok from and least tendency to dissolve iodine, and was used on all stopcocks and ground joints exposed to the halogen. (9) Commercial heptane was purified by refluxing with iodine and fractional distillation.
NOTES
111
fractive index measurements, using a Bausch and Lomb “Abbe-66” Refractometer, gave ethanol concentrations wit.h an error of less than 5 parts in 10000. Ethyl iodide/ ethanol mixture refractive indices decrease linearly with ethanol content up t,o at least 0.4 M of ethanol. Some com etitive exchanges were carried out with mixtures of metgyl, ethyl and propyl iodides. Tests with labeled ethyl iodide proved that the three halides could be separated without inducing exchange of iodine between them by fractional distillation at atmospheric pressure in the ceramic-packed column already mentioned. For the exchange experiments, samples containing equal volumes of the alkyl halides’o and the desired amount of radioactive iodine were distilled, outgassed and sealed off under vacuum as before. After several days in the thermostat, they were assayed for total, “organic” and “inorganic” activity as usual, the “organic” portion fractionated and the activities and refractive indices of successive 2-ml. aliquots determined. These measurements enabled calculation of total exchange and the proportion of the total for each halide, since the refractive indices gave the composition of each aliquot.
Results and Discussion Since in all of this work the ethyl iodide was in large excess, the appropriate equation for R, the rate of exchange, is In some instances, R was obtained from the slope of ln(1 - f) us. t plots, while in others a nomographic procedure was employed.11 With iodine in the concentration range 10-4 to 1 0 - 2 mole/l., marked inconsistencies were observed which could not be eliminated entirely. It was found that the exchange rate varied greatly with the method of ethyl iodide purification. Any one product behaved fairly consistently, but exchange half-times under comparable conditions varied in the range from 2 to 12 days a t 45” for different preparations. The most rigorously purified ethyl iodide exchanged least rapidly a.nd gave the most coherent set of data, so was deemed the “best” as alluded to above. On increasing the iodine concentration into the range 0.03 to 0.3 mole/l., and using only the “best” iodide, consistent runs under a wide variety of conditions became possible, with the results summarized herewith. (1) At 61”) using various iodine concentrations in pure ethyl iodide, the rate of exchange was found to be proportional to the first power of the iodine concentration (see Fig. 1, curve A.) The curve of best fit gives the value R = 0.09(Iz)mole/l./day. (2) Exchanges were run a t several other temperatures in the range 45 to 70”. The apparent activation energy of the process as determined from the usual linear Arrhenius plot is 19,500 cal./mole. (3) The ethyl iodide concentration was varied by dilution with heptane. With (Iz)fixed a t 6.5 X mole/l., these data resulted a t 61” (C2H51)(moles/l.) lo2 (moles/l./day)
R X
2.5 1.2
6.2 1.4
7.4 1.3
With (C2HJ) fixed a t 6.2 moles/l. and ( 1 2 ) varied, the results a t 61” were as shown in Fig. 1, curve B. It is clear that the exchange is independent of the ethyl iodide concentration in the range covered, and that it is first-order in iodine as before. How(10) The corresponding concentration values are 5.35, 4.15 and 3.41 moles/l. for methyl, ethyl and propyl iodides, respectively. (11) L. R. Darbee, J . Cham. Phys., 2 3 , 1349 (1955).
NOTES
112
B
4
?
2T-d 20 -
e. 6
Vol. 61
The mechanism of the exchange apparently involves a one-to-one ethyl iodideiiodine complex. With ethyl iodide in excess, a first-order dependence on iodine concentration is then quite reasonable.la The catalytic effect of the alcohol probably results from the rapid reduction of a fraction of the iodine t o yield iodide ions, either by the alcohol itself or by some trace impurity in it. It is well known that Iions exchange rapidly with ethyl iodide.l4 The observed increase in the exchange rate wit.h time no doubt occurs due to the further build-up of I- concentration, due to subsequent slower reduction processes.16 Financial support of this research by the U. S. Atomic Energy Commission under Contract No. AT(30-1)-1578 is gratefully acknowledged. (13) If exchange is by a n association mechanism
+ 12*
C2H61
C2HsI.Ia*
+ Is
C2H61*
ita rate will be given by
10
0
I I 0
Fig. 1.-A,
5
10 15 20 25 (12)) moles/l. X 102. exchange in pure Et1 at 61'; B, exchange in EtI-heptane mixture a t 61'.
ever, the value R = 0.20(Iz)obtained from curve B is more than double its previous magnitude. The reason for this discrepancy is unknown a t present. (4) Runs were made with ethanol contents varied up to about 0.15 mole/l., iodine being fixed at 4.25 X mole/l. and the temperature a t 61'. The rate of exchange always accelerated with time, leading to curvature in the In( 1 f ) us. t plots. But the latter were reasonably linear in the early stages of exchange, and accurate initial rates were obtainable. These rates varied directly with alcohol content, and satisfy the relation R = [0.08 O.S(EtOH)J (12) moles/l./day It is seen that the extrapolated R value for zero alcohol content is in reasonable agreement with the value obtained directly previously. (5) The relative rates of exchange of methyl, ethyl and propyl iodides were determined by the competitive procedure already described. The data were
-
+
E1ap8 ed
time (days)
(11)
3.0 17.0
0.22 .I2
% Total exchange
49.5
88.1
Relative specific activities (counts/min./ml.) Me1 Et1 PrI
8.5 6.6
1.0 1.0
0.74 .64
The mathematical analysis of competitive exchange data has been published elsewhere.12 It may be concluded that the R values for methyl, ethyl and propyl iodides are in the approximate ratio of their relative specific activities in the first experiment only, where the total exchange does not exceed 50%. (12) L. R. Darbee, F. E. Jenkins and G. 96, 603 (1956).
M. Harris, J . Chem. Phys.,
Here R is the equilibrium constant of the complex-forming reaction and the zero subscripts indicate total concentrations of the two reactants, complexed and uncomplexed. Bince ethyl iodide is in large k(Ia)o excess. it irr clear that for K of the order of only a few units, R &B found experimentally. The low value of 0.37 has been reported for K a t 25O (R. Keefer and L. Andrews, J . Am. Chem. SOC.,74, 1891 (1962)). However, their method of computation is subject t o very large errors; also, the value was obtained on the basis of an initial assumption that a negligible fraction of the iodine is complexed. This assumption is contrary to the spectral evidence, since even a dilute ethyl iodide/iodine solution in heptane shows no strong absorption peak a t 5250 A., which it would be expected to do if most of the iodine were not associated with ethyl iodide. (14) P. B. D. d e l a Mare, J . Cham. Soc., 3196 (1955). (15) A. Bagley, Trans. Faraday SOC.,94, 438 (1928).
STUDIES ON HOMOGENEOUSLY PRECIPITATED AND SUPPORTED PEROXIDES OF ZIRCONIUM, THORIUM AND OTHER POLYVALENT CATIONS BY DELBERT E. GANTZAND JACK L. LAMBERT Department of Chemistry, Kansas State Colleoe, Manhattan, Kansas Received July 66,I066
Peroxides of a number of tri- and tetravalent metal ions prepared by precipitation with internally generated hydroxide ion from the slow hydrolysis of urea in boiling water in the presence of hydrogen peroxide were studied as insoluble oxidants or adsorbant-oxidants for chromatographic-type column reactions: thorium, zirconium (hafnium), titanium, stannic and yttrium. Lanthanum peroxide was prepared by conventional precipitation with ammonium hydroxide in the presence of hydrogen peroxide after homogeneous precipitation produced a weakly reactive peroxide. The peroxides of thorium and zirconium exhibited the most desirable chemical and physical properties, and retain their activity indefinitely in dry storage. Slow crystal growth during precipitation produces discrete but somewhat gelatinous particles which are easily washed and dried at 90(1) Based upon a thesis submitted by Delbert E. Gantz in partial fulfillment of the requirements for the degree of Master of Science, Kansas State College, 1955.
