NOVEMBER 1947
933
The formation of the hydroxamate is greatly influenced by temperature. Thus, a t 110” C. the sodium hydroxamate in the reaction flask is completely destroyed on continued heating after the vessel is dry. At 55’ C. the flask may be heated for several minutes with little change in the intensity of ferric hydroxamate obtained. \Tide fluctuations in room temperature influenced the intensity of color, apparently through the cooling effect of the small quantity of reflux remaining in the flask. The effect of fluctuations in room temperature can be obviated by introducing solution A into the reaction flask without delay after the 5-second drying period or by cooling the flask in cooling tray before introducing solution A. The hydroxylamine hydrochloride and sodium hydroxide solutions may be conveniently added dropwise from a glass-stoppered dropping bottle; when stored, need not be renewed for a week. When organic or mineral acidity or alkalinity is present, the
solution must first be neutralized. The ratio of reagents is not critical; thus the volume of sodium hydroxide may be increased by 50y0even in neutral solutions, with little effect on the color formation. Ester values obtained by the above procedure on 1-mg. samples are shown in Table I compared with values obtained by the methods of the Association of Official .4gricultural Chemists ( 1 ) . The method has been successfully applied in the determination of sulfated esters. The term “ester value” used here is as defined by Jamieson ( 3 ) . LITERATURE CITED
Official Agr. Chem., Official and Tentative Methods of Analysis, 4th ed., pp. 412, 417 (1935). (2) Hill, C . T., IND.ENG.CHEM., ASAL. ED.,18, 317 (1946). (3) Jamieson, G. S., ”Vegetable Fats and Oils,” p. 340, New York, Chemical Catalog Co., 1932. (1) Assoc.
RECEIVED Sovember 15, 1946.
Determination of Thallous Chloride in Solutions of Thallous Decanesulfonate JOHN C. HENNIKER’
AND
E. C. LINGAFELTER, University of Washington, Seattle, Wash.
K’ T H E course of a n investigation of activities in solutions of Ideterminations long-chain alkanesulfonates, it was desired to make accurate of the solubility of a sparingly soluble salt in a solution containing an alkanesulfonate. Both thallous iodate and thallous chloride had been found suitable in similar studies (2, 3 ) . Since it Fas necessary for the cation to be common to both the sulfonate and the sparingly soluble salt, the problem narrowed t o the determination of iodate ion or chloride ion in the presence of a thallous alkanesulfonate. Attempts t o use the iodate were unsuccessful because of the poor end point in the iodine-sodium arsenite titration in the presence of the alkanesulfonate. For the determination of chloride-ion concentrations, the Mohr method was eliminated because a bulky yellow precipitate of thallous chromate obscured the red silver chromate and could not readily be removed. Titration t o the “clear point” was impossible because the colloidal nature of the sulfonate ion prevented precipitation of silver chloride. The Volhard method was not suitable, since the impossibility of quantitatively removing the silver chloride from these solutions caused a fading end point. Attempts were made to destroy the sulfonate ion by combustion, but such a high temperature was required that even thallous sulfate was decomposed. Attempts to destroy the sulfonate ion with various strong oxidizing agents were unsatisfactory. However, a method was evolved for precipitating silver chloride from sulfonate solutions for weighing on a semimicro scale. When the solution was 0.3 N in nitric acid, precipitation was rapid as long as there was an excess of chloride ion, but when there was as little as 0.0016 N excess silver ion, no amount of digestion gave quantitative precipitation. The method finally adopted was to add the last of the silver nitrate in increments of 5% of the total, digesting after each addition until no more precipitate was formed. Digestion a t 100°C. for one hour ensured quantitative precipitation in a readily filterable form. As a check on the method, known amounts of thallous chloride were determined in the presence of thallous decanesulfonate. Agreements t o within 1%in a solution 0.001 N in chloride were obtained. This action of the sulfonate is probably due t o the adsorption of (negatively charged) alkanesulfonate ions on the surface of the (positively charged) particles of silver chloride in the presence of excess silver ion. This adsorption would not occur on the particles in the presence of excess chloride ion. A further illustration of the protective colloid action of a long1 Present
addreas, Paraffine Companies, Inc., Emeryville 8, Calif.
Table I. Thallous Chloride in 0.02 N Thallous Decanesulfonate Solution a t 35’ C. Kormality of Thallous Chloride ( X 103) Cnsaturated Supersaturated 18.14 5.68 17.97 9.08 16.66 12.72 16.43 14.02 15.40 14.30 15.01 14.52
Days of Agitation 1 3 12
31
45 60
chain alkanesulfonate was provided by the manner in which equilibrium was approached between solid thallous chloride and an aqueous solution of thallous decanesulfonate. The solid was agitated for several weeks with sulfonate solutions both unsaturated and supersaturated with respect to thallous chloride. Data for a sulfonate concentration of 0.02 K appear in Table I. The temperature mas 35’ C. The rate of crystallization has been inhibited even more than the rate of solution. The possibility that these effects of thallous decanesulfonate might have been due to some property peculiar to this salt rather than to the sulfonate ion alone was eliminated. First, the conductivity of thallous decanesulfonate solutions was traced up to 0.05 N . The data showed a conductivity-concentration relationship similar to the calculated Onsager slope, followed by a sharp drop in equivalent conductance a t 0.030 h’,the concentration expected for the formation of micelles ( 5 ) . I n addition, the value of the limiting equivalent conductance (114.0) agrees satisfactorily with the sum of the limiting ionic conductances (117.0) of the decanesulfonate (5) and the thallous (1) ions. All these data show the thallous salt t o be similar t o other decanesulfonates that have been investigated and found t o be strong electrolytes up to a critical concentration ( 4 ) . Secondly, the possibility of slow deterioration of the thallous decanesulfonate was disproved by holding a solution of the salt a t the boiling point for a n hour. The conductivity was not changed by this treatment. LITERATURE CITED
(1) International Critical Tables, Vol. 6, p. 230, New York, McGraw(2)
Hill Publishing Co., 1926. La Mer, V. K., and Goldman, F. H., J. Am. Chem. SOC.,51,
2632
(1929)
(3) Stone, G, C. H., Ibid., 62, 572 (1940). (4) Tartar, H. V., and Wright, K. A., Ibid., 61, 539 (1939). (5) Wright, K. A , , Abbott, A. D., Siverts, V., and Tartar, H. Ibid., 6 1 , 5 4 9 (1939). RECEIVED December 11, 1946.
