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SOLUBILITY OF POTASSIUM IODIDE I N ACETONE' ROBERT LIVINGSTON
AND
RAYMOND R. HALVERSON*
School of Chemistry, Institute of Technology, University of Minnesota, Minneapolis, Minnesota Received October 6 , 10.46
The solubility of sodium iodide in acetone has been determined systematically over a range of temperatures by three independent groups of investigators (1, 6, 7), but only scattered measurements of the solubility of the other alkali iodides have been reported. In this paper are listed the solubilities of potassium iodide in acetone from +54.5" to -78.5"C. While, in general, the solubility curves of sodium and potassium iodides are similar in form, the maximum solubility of potassium iodide is about threefold less than that of sodium iodide. and it occurs a t -57.5"C.,instead of $25.7"C.as was reported for the sodium salt. The solid phase which is in equilibrium with the solution below -57.5"C. is a solvated form containing five (or six?) molecules of acetone per molecule of potassium iodide. MATERIALS
The acetone used was C.P.material and mas purified by three successive fractional distillations through a simple (all-glass) fractionating column. The final distillation was made from dry sodium iodide. ,The boiling-point range of this material was 0.1"C. Its density was 0.786 g. per milliliter at 25°C. The potassium iodide was of analytical grade, and was dried a t 125°C.before use. APPARATUS AND PROCEDURE
Equilibrium between the liquid and solid phases was obtained by stirring the mixture in an all-glass, closed vessel. This vessel was tubular, measured 3.6 cm. in internal diameter, and was 18 em. long. It was fitted with a small glass propeller-type stirrer which operated through a mercury seal. The mixture was stirred for from 1 to 6 hr. (after attaining thermal equilibrium), the longer times being used a t the lower temperatures. Each mixture consisted of about 70 ml. of solvent and sufficient potassium iodide to ensure an excess of a t least 10 g. When the samples were to be withdrawn, the stopper supporting the stirrer was replaced by a second ground-glass stopper, provided with a 3-mm. I.D. sampling tube and an air inlet. This stopper and sampling tube were brought to the temperature of the thermostat before being inserted into the reaction vessel. The sampling tube extended t o within 1 cm. of the bottom of the reaction vessel and terminated in a small fritted-glass filter. By applying (dry) The majority of the experimental measurements and computations reported here were performed by R. R. Halverson during the winter of 1941, while he was a n undergraduate student at the University of Minnesota. Preparation of the work for publication has been delayed by waRtime conditions. * Deceased September 1, 1945. 1
It
2
ROBERT LIVINGSTON AND RAYMOND R. HALVERSON
air' pressure to the inlet tube, approximately 25 ml. of the solution was forced directly into a specially designed weighing bottle. After a lapse of 3 or 1 hr., a second sample was withdrawn to determine whether equilibrium had been attained. The weighing bottles were of approximately 50-ml. capacity and mere completely closed except for two short, vertical, 8-mm. tubes. These tubes were closed with ground-glass caps. After the solution had been forced into the bottle, the caps were replaced, and the bottle was weighed to the nearest 0.05 g. The solvent was then evaporated by a current of dry air, the bottles being kept near the boiling point of acetone on a hot plate. The weight of the dry salt was determined by weighing the bottle with an analytical balance, the usual precautions being maintained. Temperatures above 4°C. were maintained with a water thermostat, which was provided with an adequate stirrer and a sensitive regulator. At temperaTABLE 1 Summary. of. solubility measuremknts TEMPERATURE (CORRECTED)
t
DURATION OF STIRRING
WEIGHT OF SAXPLE
SOLUBILITY
"C.
hours
grams
weight per cent
2.8 1.5 2.0 2.5 2.3 3.3 4.0 5.8 6.3 6.3 6.0 4.5 5.3
17.8 25.0 26.2 27.0 27.9 28.7 27.6 21.8 25.8 28.3 31 .O 32.5 29.1
0.84 0.89 1.30 1.330 2.02 3.87 6.08 9.2 9.79 9.79 9.0 6.84 4.15
54.50 48.75 26.54 25.00 4.62 -26.3 -44.8 -56.4 -57.0 -57.6 -58.5 -64.4 -78.5
,
tures below 0°C. a dry-ice cooled thermostat of the type discussed by G. B. Heisig (2) was used. The temperature of -78.5"C. was obtained by surrounding the reaction vessel with dry ice. The dry ice .was contained in a Dewar flask, and air was displaced from it with a current of carbon dioxide. Special experiments were performed to determine the time required for the solution to attain the temperature of the thermostat. In the upper range the temperature was measured with a calibrated 0.1"C. mercury thermometer. Temperatures in the lower range were recorded with a toluene thermometer, taking care to avoid hysteresis effects. This thermometer was carefully calibrated a t the freezing points of pure water, mercury, and chloroform and at the sublimation point of carbon dioxide. THE SOLUBILITY MEASUREMENTS
The results of the solubility measurements are summarized in table 1. The uncertainty of the tabulated temperatures and solubilities is probably less than f5 in the last place. The temperatures in the upper range were controlled and
3
SOLUBILITY OF POTASSIUM IODIDE IN ACETONE
measured within f0.05"C.; in the low temperature range, within 0.5"C. Estimates of the accuracy of the solubility measurements are based largely upon the reproducibility of successive determinations of the same mixture. With few exceptions, the difference between two successive determinations of the solubility was less than 1 per cent df the listed value, but nearly always represented an increase with increasing time of stirring. At 25"C., samples taken a t 1, 14, and 23 hr. yielded solubilities of 1.318, 1.329, and 1.329 per cent, respectively. At -56.4"C. only one measurement was made, the value being 1.23 per cent. At -58.5"C. values of 8.58 and 8.98 per cent were obtained at 5.5 and 6.0 hr., respectively. Preliminary values of 8.26 and 8.56 per cent, which.were obtained, respectively, a t 3.5 and 4.3 hr. a t -58.5"C., have been excluded from the table. TABLE 2 Summary of measurements of the degree of solvation
. . . ..
.I
Weight of sample (grams). . . . . . . . . . Number of acetone molecules.... . . . . . .
3.1 5.1
I :::I
0.9 5.8
1
2.5 5.9
1
2.0 4.9
DETERMINATION OF THE EMPIRICAL FORMULA OF THE SOLVATE
Several attempts were made t o determine the empirical formula of the solid sohate. The solid was prepared by saturating a solution at -58°C. and precipitating the solvate by chilling the solution to -78°C. A sample of the solid which separated was transferred with a spatula t o a pad of filter paper which had been chilled with dry ice, The sample was worked on the pad with the spatula, or squeezed in a thin layer between two pads, for about 1 min. in an attempt to remove the mother liquor. The sample was then quickly transferred t o a weighing bottle and weighed. After driving off the solvent, the sample was again weighed. Table 2 summarizes the results obtained. Since it was uncertain how these results were affected by the presence of mother liquor adhering t o the crystals or possibly by incipient decomposition of the solvate, a check experiment was performed using a power-driven centrifuge t o dry the crystals. In this experiment the solvate, prepared as before, was transferred with a spatula to a chilled, sintered-glass crucible. The crucible was supported in the upper part of a brass centrifuge cup by means of an adapter. The remainder of the cup was packed with dry ice. The centrifuge was run for about 3 min., a t the end of which time the cup still contained about half of the dry ice originally present. The crucible was then transferred to a weighing bottle, and the weight of the salt determined before and after driving off the acetone. The corresponding number of acetone molecules per molecule of potassium iodide was 4.97: Although it cannot be claimed that these measurements establish it with certainty, it is probable that the formula for the solvate is KI * 5C3HsO. DISCUSSION
The solubilities of sodium and potassium iodides, as well as scattered data for the other alkali iodides, are plotted as a function of temperature in figure 1.
4
ROBERT LIVINGSTON AND RAYMOND R. HALVERSON
A semilog scale is used in the graph t o facilitate comparison of the data, since the solubilities vary by three hundred fold. The data of table 1 are represented by circles. They are in excellent agreement with the two determinations of Walden (7), which are indicated by solid dots. The older data of von Lasczynski (4)(plotted as X's) and the two values of Lannung (3) (plotted as crosses) shorn
1
I
\cs 0.1
1
SOLUBILITP OF POTASSIUM IODIDE I N ACETONE
5
maximum solubility of sodium iodide is threefold greater and corresponds to a temperature 83°C. higher than the potassium iodide maximum. The solubility of the unsolvated sodium iodide decreases much less rapidly with increasing temperature than does that of the unsolvated potassium iodide. It is possible, by plotting log x vs 1/T and taking the slopes of the curves a t the transition temperature, to compute approximate values for the heats of solution of the solvated and unsolvated salts and (by difference) the heat of solvation of the solid. Applying this method t o the sodium iodide data (6) we obtain NaI (s)
+ 3C3He0 (1) = NaI.3(C3HsO) (s);
AH = -9,500 cal.
It is difficult to obtain a reasonably precise value for potassium iodide, since the slopes of the curves change rapidly as they approach the transition point. However, the data do indicate that the heat of the solvation reaction, K I (s)
+ nC3HeO (l)= KI.nC3HsO (s)
lies between - 14,600 and - 18,000 cal. Examination of analogous data (1) for alcoholates of a variety of salts indicates that the ratio of the heat of solvation to the number of solvent molecules attached is independent of n and Me in the formula MeX.n(alcoho1). If we assume this generalization to hold in the present case, we can compute the heats of solvation corresponding to n = 5 and to n = 6 as 5/3(-9,500 cal.) = -15,800 cal. and 6/3(-9,500 cal.) = - 19,000 cal., respectively. Either of these values is within the range obtained directly from the potassium iodide data. The curves for lithium and cesium iodides are based on the data of Lannung (3), while that for rubidium iodide is from Walden's data (8). Lannung's values (3)3 for rubidium iodide, which are not represented on the graph, show the same temperature dependence but are about 20 per cent lower. While the solubility data for these salts are very meager, they are consistent with the more complete data for sodium and potassium iodides in suggesting that the solubilities of the alkali halides form a family of related curves. SUMMARY
The solubility of potassium iodide in acetone increases from 4.15 per cent (by weight) a t -78.5"C. to 9.80 per cent at -57.5"C.; it then decreases as the temperature is further increased, being 1.330 per cent at 25.0"C. and 0.84 per cent a t 54.5"C. Below -57.5"C. the stable form is a solvate, containing five (or possibly six) molecules of acetone per molecule of potassium iodide. REFERENCES (1) BELL,W. R. G . , ROWLANDS, C. B., BAMFORD, I. J., AND JONES,W. J.: J. Chem. SOC. 1930, 1927. These values are incorrectly listed in Seidell's Solubilities of Inorganic Compounds (1940) as being measured at 18" and 25°C. instead of 18" and 87°C.
6
JOHN A . BISHOP
(2) HEISIG,G. B . : Ind. Eng. Chem., Anal. Ed. 8, 149 (1936). (3) LANNUNG, A.: Z . physik. Chem. A161, 265 (1932). (4) LASCZYNSKI, 8. VON:Ber. 27, 2285 (1894). (5) LOYD,E., BROWN,C. B., GLYNWYN, D., BONNELL, R., AND JONES,W. J.: J. Chem. SOC. 1928, 658. (6) MACY,R., AND THOMAS, E. W.: J . Am. Chem. SOC. 48, 1547 (1926). (7) WADSWORTH, A. E., AND DAWSON, H. M.: J. Chem. SOC.129, 2784 (1926). (8) WALDEN,P . T.: Z. physik. Chem. 66, 712 (1902).
COVALENT ADSORPTION ON BASE-EXCHANGE RESINS. I
THEADSORPTION OF MONOBASIC ACIDS JOHN A. BISHOP' Chemical Laboratory, University of Delaware, Newark, Delaware Received July 26, 19.@
Since the pioneering work of Adams and Holmes (1) on the use of resins as base exchangers, a large number of patents have been obtained describing resinous and carbonaceous base-exchanging materials. On the theoretical side of this type of adsorption, the first attempts t o investigate the phenomena were made by Bhatnagar and coworkers (3, 4) and by Broughton (2, 5 ) . Walton (9, 10) and Nachod (6, 7) have discussed and summarized the results on cation exchangers. Comparatively little investigation has been made of the adsorption of anions except chloride and sulfate (2, 6, S ) , aside from the early work of Bhatnagar. The problem may be considered as one involving the formation of covalent bonds between the nitrogen of the resin and the hydrated proton of the acid, the anion being retained by electrostatic attraction as in the neutralization of ammonia by an acid. EXPERIMENTAL
A . Materials The adsorbent used in this series was Amberlite IR-3, obtained from the Resinous Products and Chemical Company through the courtesy of Dr. F. J. Myers. It was sized by sieving through No. 20 on No. 40 U. S. Standard sieves. It was then freed from soluble impurities by extraction with water, using a Soxhlet extractor, until a colorless extract was obtained. The residue was dried a t 110°C. and stored over calcium chloride. This treatment resulted in a material which did not discolor distilled water. The acids used mere either Baker's C.P. grade or from the Eastman Kodak Company. 1
Present address: Moravian College for Women, Bethlehem, Pennsylvania.
