E. Roig, I. G. Rieckehoff, C. 5. Russo, and J. D. Curet1
Puerto Rico Nuclear Center University of Puerto Rico Rio Piedras, P. R.
Radioisotope Demonstration of Common Ion Effect 0n Solubility
T h e commercial availability of many radioisotopes at microcurie levels which do not require that the purchaser possess an AEC l i c e n ~ ehas , ~ placed new tools in the hands of chemistry teachers. Radioactive atoms behave chemically in essentially the same way as their stable isotopes, and their radiations are easily detectable; these facts make them excellent means for demonstrating chemical principles. Often a chemical analysis can be replaced by a simple radioactivity measurement with consequent saving of time and reagents. I n other cases, it is possible to show better certain effects which are not so obvious when illustrated by conventional methods. The calculations necessary for the interpretation of results usually are very simple and straightforward. I t is not necessary for the person who does the demonstration to be an expert in the field of radiochemistry. I t is sufficient that he be acquainted with the basic techniques and precautions for handling radioisotopes. This paper reports a direct demonstration of the effect of common ion concentration on the solubility of a uni-univalent salt, thallium(1) chloride. Thallium@) chloride has been extensively used for solubility studies since the beginning of the cent,ury ( 1 , 2 ) and can be labeled either with chlorine-36 or with a thallium isotope. Thallium-204 was used here. The Experiment Thallium-204 can be obtained as thallium(1) nitrate in HSOs solution from Oak Ridge National Laboratory. The acidity of the solution is below 3N and the concentr&m of the activity not less than 1 me/ml. Thallium-204 decays with a half life of 4.0 years by bela emission (maximum energy 0.765 Mev) and by electron capture (2'7J. Thallium(1) nitrate solution, 0.10A4, is labeled by adding enough TI-204 solution to 100 ml of O.lO.M TINOI so that the resulting solution has an activity from 1 to 5 microcuries per ml. For the solubility study, the following solutions are convenient: Solution 1,0.60M KNOl Solution 2,0.05Min TINOs and 0.55M in KNOs Solution 3,O.lOMin TINOXand 0.50M in KNOB Solution 4,0.05M in KC1 and 0.55M in KNOI Solution 5,0.10M in Kcland 0.50Min KNOs
The day before the demonstration is to he given, prepare the precipitate of thallium(1) chloride by placing in each of five conical 15-ml centrifuge tubes, 2 ml of 0.10M thallium@) nitrate labeled with thallium-204 and 2 ml of 0.20M KC1 solution. Let stand overnight. On the following day, centrifuge the tubes, discard the supernatant, drain each tube well, and wash each Present address: Desn, Faculty of Natural Sciences, University of Puerto Rico. 2 Nucleonic Corporation of America., 196 DeGraw St., Brooklyn 31, N. Y., supplies license-free amounts of several radioisotopes.
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Journal of Chemicol Education
precipitate with 1 ml of ~ m t e r . Centrifuge again, decant the supernatant, drain well. Add to first tube 1 ml of solution 1, to the second tube 1 ml of solution 2, to the third tube 1 ml of solution 3, to the fourth tube 1 ml of solution 4, and to the fifth tube 1 ml of solution 5. Shake each tube well after addition of the solutions and continue shaking the tubes in a regular pattern for a t least 10 minutes. (It is better to use a shaking machine but it is not indispensable for semiquantitative results.) Centrifuge the tubes and measure an aliquot (50-200 h ) from each of the supernatants into polystyrene bases mounted on cardboard or aluminum. Dry under an infrared lamp, cover samples with cellophane, and count them in an end-window Geiger counter or a thin-window flow counter. Precipitates are allowed to stand overnight in order to decrease the rate of isotopic exchange which is known to be lower for aged precipitates (3). The rate of isotopic exchange between the labeled solid thallium(1) chloride and thallium(1) ion in solution was found to be insignificant under -the conditions of these experiments. Only 10 minutes are allowed for the precipitates to be in contact with the solutions wit,h occasional stirring. These condit,ions were found to be satisfactory for saturation to be estahlished yet to keep the time of the demonstration as short as possible. I t is evident however, that for more accurat,e results, not only a shaking apparatus but also a longer agitation period and a constant temperature bath are required. Also, solutions with common ion concentrations below 0.05M should be studied in order to include a larger number of points in the graphical treatment of the data. The samples are dried before counting in order to minimize self-absorption effects. This t.akes from 5 to 10 minutes which might be used by the t,eacher t,o explain the operation of the counter and to obtain the background counting rate. The samples should contain enough activity so that counting for a short time,,, (1 or 2 minutes) will give sufficient counts for gqod stat,istical significance. . r +. After the samples are counted, the results can he tabulated on the blackboard and discussed. The explanations for the results obtained may be made as simple or rigorous as the level of student training permits. A relatively high ionic strength, 0.6M, was chosen in order to neglect the contribution of the ions coming from the dissolution of the precipitate. Potassium nitrate was used as the inert salt because of its availability. However, since nitrate is known to associate somewhat with thallium(1) ion ( d ) , perchlorate mould be a better inert anion.
Some typical results obtaiued by the use of the method are presented in Table 1. I n each case, aliquots of 0.10 ml were counted in a thin window flax proportional count,er. Table 1 .
Relative Solubilities of TIC1 in the Presence of CI- Ion (Constant I o n i c Strength = 0.60)
Solution 0 60M KNOI 0.05M TINOI f 0 55M KSOa 0 10M TlNOs 0.50M KXOa 0.05M KC1 0.55M KXOI 0.10M KC1 0.50M KXOa
+ + +
Piumher Arithmetic Per cent of mean average (cpm) deviation samples 4 13415 2 6 6329 2 5 4508 4 6 6654 3 6 4057 3
When TI- is the common ion, assuming no exchange with the precipitate, K = (A/a z)(A/a) or A2 = Ka2 - aAz. I n both cases, a plot of A2 versus Ay or AZ should be a straight line with intercept K a Z and slope -a. Therefore, K can be calculated from this value, since K = intercept/slope2. Figure 2 represents this treatment as applied to the experimental data. The chloride data falls excellently on a straight line and the thallium(1) data agrees with the same line as expected but shows more scatter. The line has an intercept of 1.80 X lo8 and a slope of 4.1 X 105giving a value of 1.07 i 0.2 X 10W8 for K,the solubility product. The error assigned to K represents the extremes of the tballium(1) data. From this value, the solubility of TIC1 at room temperature (approximately 2S°C) is found to be 0.033 moles per liter. No data is available for the solubility of TIC1 under the conditions used in this experiment. By making certain rough interpolations using data from Seidell (6) the solubility of TlCl under these conditions was estimated as 0.030 mole per liter. This value shows a good agreement xith the experimental value considering the nonrigorous control of conditions in a demonstration experiment.
+
I
I
om
aoa
I
o
TI' MOL~RTY
005
010
CI- MOL~RITY
Figure 1. Relotive d u b i l i t i e r of tholliumlll chloride in solutions of KC1 a n d T I N 0 2 with constant ionic strength.
Figure 1 shows the effect of increased commou ion concentration on the solubility of TlCl. The similarity of the curves represented in Figure 1 is evidence that the common ion effect in depressing the solubility of mi-univalent salt,s is of the same order of magnitude for hot,h the eation and anion, provided other equilibria are eit,her absent or insignificant under the conditions considered. The data can be interpreted quantitatively to show that it conforms to a simple equilibrium law and can be used to determine the value of Ksp. = (TI+) (CI-). Although it is known that additional equilibria such as TI+ C1- d TIC1 solution and TIC1 Cl- d T l C k exist (5),they can be neglected under the conditions of t,he experiment. At ionic strength = 0.60,the equilibrium constant of the first has a value of approximately 1 and t.he second is still smaller. Therefore, a t (CI-) = 0.10M (the highest chloride concentration used in the experiment) only about 10yo of t,he thallium in solution is present as undissociated TlC1.
+
Let A
+
=
a =
y
=
r =
measured activity of an aliquot of the supernatant (same volume aliquot in every case) specific activity of solid in units such that A/a is the concentration of the T1+ which hae appeared in the solution by dissolut,ion of the solid concentration of added C1concentration of added TI+
When CI- is the common ion, K or A2 = K a 2 - aAy.
=
(A/a) (A/a
+ y)
Figure 2.
Grophicol determination of solubility product.
We are indebted to Prof. R. W. Dodson, Chairmau of the Chemistry Department of Brookhaven National Laboratory, for his suggestions concerning the presentation and the interpretation of the data. Literature Cited (1) BRAY,WILLIAM C., AND WINNINGHOFF, W. J.,J. Am. Chem. Soc., 33,1663 (1911). (2) BRAY,WILLIAMC., J. Am. Chem. Soc., 33,1673(1911). I. M., AND ROSENBLUM, CHARLES,J . A m . (3) KOLTHOPF, Chem. Soe.. 56.1658 11934). , (4) NAIR,V. s.,'A& NANCOLLAS, G. H., J . Chem. Soc., 1957, 318. (5) Hu, K. H., AND Scow, A. B., J . Am. Chem. Soc., 77, 1380 (1955). (6) SEIDELL,A., "Solubilities of Inorganic and Metal Organic Compounds," 3rd ed., D. van Nostrand Co., Inc., S e w York, 1940, Vol. 1, p. 1547.
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Volume 38, Number 7, July 1961
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