MAY 15, 1936
ANALYTICAL EDITION
assumption was made that each mole of peroxide reacted with two equivalents of ferrous sulfate. The comparison of the two methods is presented in Table I, the results being reported in terms of gram equivalents of active oxygen per liter of hyclrocarbon. The lower results given by the method of Yule and Wilson are probably due to the incomplete reduction of the peroxides, for the reacting substances are, on the whole, concentrated in two different layers. The more peroxide there is present, the more difficult it becomes to obtain a quantitative reduction. Yule and R7ilson have noted that an increase in peroxide concentration does not give a proportionate increase in the quantity of peroxide detected by their method.
199
Summary
-4method is presented for the colorimetric estimation of peroxides in unsaturated organic compounds, based on the oxidation of ferrous sulfate in the oresence of ammonium thiocyanate, using absolute methyl dcohol as the solvent.
f
Literature Cited (1) Clover a n d H o u g h t o n , Am. Chem. J . , 32, 55 (1904). (2) M a r k s a n d Morrell, Analgst, 54, 503 (1929). (3) Yule a n d Wilson, IKD. ENG.CHEM.,23, 1254 (1931). RECEIVED January 22, 1936.
Volumetric Determination of Iodides bv Ceric Sulfate 0
J
An Application of the Indicator o-Phenanthroline Ferrous Ion DAVID LEWIS, College of the City of New York, New York, N. Y.
T
HE direct oxidative titration of iodides to a visual end
point depends, in most procedures, upon the conversion of the initially liberated iodine to the cationic or covalent state. The end point is the disappearance of iodine, and it is in the manner of determining this disappearance that the various methods differ. Each is subject to certain disadvantages, either of manipulatibn or of estimating the end
point. In Andrews’ method (1) an immiscible organic solvent, such as chloroform, must be used t o detect the oxidation of thetiodine t o iodine monochloride. Addition of oxidizing agent, with shaking after each addition, is continued until the violet iodine color is discharged from the layer of inert solvent. Lang’s procedure ( 6 ) is more direct. Titration is continued t o the disappearance of the blue starch-iodide color which occurs at the Conversion of the iodine to iodilie cyanide. To minimize the danger of working with an acid solution of a cyanide, it is necessary to use long, narrow-necked flasks. In Berg’s method (a) the iodine liberated in the oxidation reacts with acetone to form iodoacetone. Starch is the indicator. As the end point is approached each drop of oxidizing agent produces a blue color which is slowly discharged. The end point is reached when further addition of oxidant no longer produces a blue color, The most common oxidizing agent is potassium iodate, although Swift i(7) has shown that potassium permanganate, potassium dichromate, and ceric sulfate can be used in the Andrews procedure, and Berg (3) has shown the utility of potaasium bromate in the cyanide method. I n the method to be described the disadvantages of these methods are eliminated by avoiding the use of an iodine end point. A previous attempt in this direction was made by Hahn and his oo-workers (5),who proposed titration to a permanganate end point, without, however, eliminating the inconvenience of a two-phase system. The iodine liberated in the oxidation had to be extracted (with ethyl acetate) to permit detection of the end point. The present method is based upon Berg’s procedure for the elimination of iodine by the acid catalyzed iodination of acetone. By the use of ceric ion as oxidizing agent and o-phenanthroline ferrous ion as indicator, the titration of iodides can be performed rapidly, precisely, and accurately.
Reagents Two ceric sulfate solutions were prepared by dissolving ceric ammoriium sulfate in M sulfuric acid, and were standardized
against sodium oxalate by the method of Walden, Hammett, and Chapman (8). Solution I was 0.1074 M ; solution 11, 0.09982 M . An approximately 0.1 M solution of purified potassium iodide was used throughout these experiments. It was standardized against the ceric sulfate solutions, the end point being determined electrometrically-a procedure shown to be exact by Willard and Young (9). Acetone was of reagent grade and the potassium bromide, sodium chloride, and sulfuric acid were of c. P. grade. Blanks on these materials in the amounts used in the experiments required a fraction of a, dr8p of ceric sulfate to change the color of the indicator. The solurnon of o-phenanthroline ferrous sulfate was 0.025 M .
Method of Analysis A measured volume of the iodide solution is treated with 25 ml. of acetone, 10 ml. of 9 M sulfuric acid, and water to make the volume 100 ml. After adding one drop of o-phenanthroline ferrous sulfate solution, the mixture is titrated with the ceric sulfate until the pink color of the indicator changes to a pale blue. The end point is sharp and lasts several minutes. The rate a t which the oxidant is added does not affect the results. At the start of a titration rapid addition of the ceric sulfate may cause the solution to be colored brown by free iodine, which rapidly disappears on interrupting the titration and stirring a few seconds. It is desirable to conduct the titration in flasks, since iodoacetone is a lachrymator.
Results Table I illustrates the precision with which this titration can be performed. There is no difficulty in obtaining checks better than one part per thousand. From these results the normality of the iodide solution is found to be exactly twice the molarity as determined potentiometrically. Assuming monoiodination of the acetone, the reaction may be represented as KI
+ 2Ce(SO4)2 + CsH60 = KHS04 + Cee(SOa)s 4- C&OI
It is unnecessary to adhere strictly to the conditions for the titration as given above. Equally good results are obtained if 5 or 15 ml. of acid are used or if the solution is diluted. Such changes merely affect the rate of reaction of the iodine
INDUSTRIAL AND EKGINEERING CHEMISTRY
200
TABLE I. PRECISION ATTAINABLE Potassium Iodide MI.
Ceric Sulfate
M1. Solution I, 0.1074 M 18.67 18.69 18.70 18.69 18.70 Av. 18.69 Av. deviation 0.8 part per 1000
10 I O 0
Solution 11, 0.09982 M 10.00
20.10 20.10 40.20 10.04
20.00 5.00
Av. deviation 0.1 part per 1000
TABLE11. EFFECTOF VARYINQ CONCENTRATION OF ACIDAND VOLUMEOF SOLUTION (10.00 ml. of potas ‘um iodide and 25 ml. of acetone in each titration. sulfate required, 18.69 ml.) Volume a t Start 9 M H&Od Ceric Sulfate Used
%kit
.M2.
MI.
.M1.
100 100 100 100 145 100 200
1 5 15 20 30 30 20
18.75 18.70 18.68 18.61 18.60 18.43 18.6Y
TABLE111. EFFECTOF BROMIDE (10.00 ml. of potassium iodide and 25 ml. of acetone in each titration. Ceric sulfate required, 18.69 ml.) 0.1 .MKRr Volume a t Start 9 .M H P S O ~ Ceric Sulfate Used
‘Ml
.
MI.
MZ.
M1.
VOL. 8, NO. 3
low results and a fleeting end point. The data in Table I1 show the effect of varying these conditions. EFFECTOF BROMIDE.Measured volumes of 0.1 M potassium bromide were added to the iodide and the titration was performed as described above. I n general, more ceric sulfate is required than for iodide alone. The larger the ratio of bromide to iodide, the greater is the excess of ceric sulfate necessary to reach an end point lasting 1 minute. This interference can be almost entirely eliminated by diluting the solution sufficiently while keeping the acid concentration approximately constant. Within the limits investigated it was possible to titrate the iodide with an accuracy of 3 to 4 parts per thousand when the bromide-iodide ratio was 5 to 1. The results of several of these titrations are given in Table 111. EFFECTOF CHLORIDE.Moderate amounts of neutral chloride do not interfere. Titrations in the presence of 0.5, 1.0, and 5.0 grams of sodium chloride required 20.10, 20.16, and 20.23 ml. of ceric sulfate, compared to 20.10 ml. for iodide alone. I n the titration with 5.0 grams of chloride present, precipitation of the salt occurred. The excess ceric sulfate required with the larger amounts of chloride may be due to the formation of hydrochloric acid, since titrations of iodide alone in which hydrochloric acid was substituted for the sulfuric acid used excess ceric sulfate.
Summary In the presence of acetone and sulfuric acid iodides may be titrated quantitatively with ceric sulfate to a visual end point, using o-phenanthroline ferrous ion as indicator. The effect of bromides and chlorides on this titration has been determined.
Literature Cited (1) Andrew,
a 35
ml. of acetone.
Kith the acetone, which is proportional to the concentration of acid and of acetone (4). Too little acid is undesirable, since it leads to slightly high results and the titration is timeconsuming, while too high a concentration of acid leads to
LYLE 0. HILL Central Y. M,C. A. College, Chicago, Ill.
T
HE method of standardizing sodium thiosulfate against
copper and of determining copper in samples which are soluble in nitric acid as outlined by Gooch and Heath (1) requires long evaporations. Kendall (8) has pointed out that these evaporations eliminate nitrous acid. The use of urea for the elimination of nitrous acid was worked out independently by the author. Afterwards, in a careful study of the literature it was found that Koelsch (3) suggested a procedure essentially the same as that given below and that Pozzi-Escot (4) also suggested the use of urea, employing, however, a much longer procedure. It is felt that this simple method has been overlooked and should be called to the attention of the analytical chemist. Results with this method agree with those of Gooch and Heath (1) within one part per thousand.
L. W., J.Am. Chem. Soc., 25, 756 (1903).
(2) Berg, R., 2. anal. Chem., 69, 369 (1926). (3) Ibid., 69, 1 (1926). (4) Dawson, H. M., J . Chem. SOC.,1927, 458, (5) Hahn, F. L., and Wolf, H., Chern.-Ztg., 50, 674 (1926); Hahn, F. L., and Weiler, G., Z. anal. Chem., 69, 417 (1926). (6) Lang, R . , 2. anorg. a2Zgem. Chem., 122, 332 (1922); 142, 229 (1925); 144, 75 (1925). (7) Swift, E. H., J . Am. Chem. SOC.,52, 894 (1930). (8) Walden, G. H., Hammett, L. P., and Chapman, R. P., Ibid., 55, 2649 (1933). (9) Willard, H. H., and Young, P.,Ibid., 50,1368 (1928). RECEIVED February 13, 1936.
Procedure Weigh a sample of pure copper, 0.2 to 0.3 gram. Dissolve the sample in 2 t o 5 ml. of concentrated nitric acid. Add 0.5 ram of urea and heat to boiling. Cool, adjust the acidity by afding 6 N ammonium hydroxide until a white precipitate is formed, dissolve the precipitate with 6 N acetic acid, and add 5 ml. in excess. Add 3 grams of potassium iodide, allow to stand 2 minutes, and titrate the liberated iodine with sodium thiosulfate solution. Starch is used as an indicator ( 2 ml. of l per cent solution) and should be added about 1 ml. before the end point is reached. Variations in concentration of KO3- and NH4+ have no effect on the final precision.
Literature Cited (1) (2) (3) (4)
Gooch and Heath, 2. anorg. Chem., 55, 199 (1907). Kendall, J. Am. Chem. Soc., 33, 1947 (1911). Koelsch, Chem.-Ztg., 37, 753 (1913). Poaai-Escot, E., Ann. chim. anal., 18, 219 (1913).
RECEIVED December 23, 1935.
