Colorimetric Estimation of Peroxides in Unsaturated Compounds CHARLES A. YOUNG, R. R. VOGT, AND J. A. NIEUWLAND, University of Notre Dame, Notre Dame, Ind.
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HE chief methods of estimating the peroxides in unsaturated organic compounds are based on the oxidation of either potassium iodide or a ferrous salt by the peroxides. For example, Marks and Morrell (2) use a solution of potassium iodide in glacial acetic acid. After the iodine is liberated, the solution is diluted with water and the free iodine titrated with standard sodium thiosulfate. A considerable error may be introduced by the addition of iodine to the unsaturated linkage, particularly since the rate of addition in many cases is greatly accelerated by the presence of peroxides. Yule and Wilson (3) estimate peroxides in cracked gasoline by shaking it with an acidified solution of ferrous sulfate and ammonium thiocyanate in a 50 per cent acetone-water mixture. The ferric thiocyanate which is formed is titrated with a standard titanous chloride solution. Several improvements in this method are desirable. Instead of using as the solvent 50 per cent acetone, in which most hydrocarbons have a limited solubility, it is preferable to use a solvent which is miscible with hydrocarbons. It is also desirable to estimate the ferric thiocyanate formed by the peroxides by some means other than standard titanous chloride, in order to avoid the difficulties inherent in the use of this reagent. Since minute quantities of ferric thiocyanate produce an intense coloration when in solution, it is possible and advantageous to determine the ferric thiocyanate colorimetrically. The convenience and rapidity of colorimetric procedures are well known. In selecting a solvent there are two important factors to be considered : (1) the solvent should dissolve an appreciable quantity of the inorganic salts used, and (2) the solvent should be miscible with the organic compounds which are to be analyzed. Absolute methyl alcohol meets these requirements.
Preparation of the Reagent The reagent is prepared by dissolving 5 grams of ammonium thiocyanate and 5 cc. of 6 N sulfuric acid in 1000 cc. of absolute methyl alcohol, and then saturating the solution with pure ferrous ammonium sulfate. Shaking the solution with the finely pulverized salt for a few minutes is sufficient. The faint pink color which is formed may be evaluated by comparing it with a suitable color standard. The results of analyses may then be corrected for the trace of ferric thiocyanate present in the reagent. The faint pink color in the reagent does not deepen appreciably over a period of an hour or so. The usefulness of the reagent may be extended to much longer periods by keeping the reagent in an inert atmosphere.
methyl alcohol. For colorimeters of the Duboscq type, it is recommended that the concentrations of the ferric chloride .solutions range from 0.00004 to 0.001 mole per liter. Three concentrations are sufficient for ordinary purposes. In case any ferric chloride has been reduced to the ferrous state, a small amount of actiGe peroxide may be added.
Procedure If the compound being analyzed is a liquid having a moderate peroxide concentration, it is added directly to the reagent by means of a small pipet or buret. Otherwise a solution in absolute methyl alcohol is made, and the concentration so adjusted that, when quantities of the solution ranging from 0.05 to 0.5 cc. are added to 10-cc. portions of the reagent, a color equivalent to 0.00002 to 0.0002 mole of ferric thiocyanate per liter is produced. With many compounds, such as butylacetylene and 1-hexene,the color reaches a maximum intensity within a few seconds, and the solution is then immediately compared with a color standard having approximately the same concentration. Some peroxides, such as the one found in diamylene, react with the reagent rather slowly. In such cases the reaction may be greatly accelerated by heating the solution to just below boiling for 4 or 5 minutes. The quantity of peroxide detected by this method is pro or tional to the amount present in the solution. The results o f an analysis are reproducible. ,
Discussion Potassium iodide is more easily oxidized than is ferrous sulfate, as the molal electrode potentials of the ferrous and the iodide ions show. Certain peroxides, such as benzoyl peroxide, readily liberate iodine from potassium iodide, but react extremely slowly, if a t all, with ferrous sulfate. Consequently, if some organic material contains a wide variety of peroxides, it is possible that a portion of the peroxides is not determined by a method using a ferrous salt as the reducing agent. Yule and Wilson (3) found that after an exhaustive treatment of cracked gasoline with ferrous sulfate, there still remained a small amouht of peroxide which oxidized potassium iodide. Severtheless, since potassium iodide cannot be used with many peroxides of unsaturated compounds, the ferrous sulfate methods are very useful for comparative purposes even where complete reduction of all peroxides is not obtained. TABLEI. COMPARISON OF METHODS Yule and Wilson Cracked gasoline Butylaoetylene Diamylene
0.025
0.23 0.30
Colorimetric 0.036 0.29 0.54
The accuracy of the colorimetric method presented here has been studied by means of dilute solutions of hydrogen peroxide of known concentrations and by means of pure preparations of succinyl peroxide (HOOC.CHz.CHz.C00)2. For example, the concentration of a hydrogen peroxide solution as given by the standard potassium permanganate method was 0.0279 mole per liter, while the colorimetric procedure gave 0.028 mole per liter. The succinyl peroxide was prepared by the method of Clover and Houghton ( I ) , and was recrystallized twice from acetone. The calculated weight per cent of active oxygen in succinyl peroxide is 6.83 per cent; obtained 7.0 per cent. For fuller evaluation of the merits of the colorimetric procedure, it was compared with the method of Yule and Wilson (3). In caIculating the results of an analysis, the reasonable
Preparation of the Color Standard Solutions of ferric thiocyanate in absolute methyl alcohol are used as color standards. In order to make the colorimetric comparisons conveniently and accurately, the concentration of the standard solution should not differ more than two- or threefold from that of the solution whose color is being determined. Furthermore, since solutions of ferric thiocyanate in methyl alcohol fade slowly on standing, the standard solutions should be freshly prepared each day. This may be accomplished by adding ammonium thiocyanate and sulfuric acid, in the same proportions used in preparing the reagent, to standard solutions of ferric chloride in absolute 198
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.
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Literature Cited (1) Clover and Houghton, Am. Chem. J . , 32, 55 (1904). (2) Marks and Morrell, Analgst, 54, 503 (1929). (3) Yule and Wilson, IKD. ENG.CHEM.,23, 1254 (1931). RECEIVED January 22, 1936.
Volumetric Determination of Iodides bv Ceric Sulfate 0
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An Application of the Indicator o-Phenanthroline Ferrous Ion DAVID LEWIS, College of the City of New York, New York, N. Y.
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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
