V O L U M E 20, N O . 11, N O V E M B E R 1 9 4 8 cedure is therefore not recommended where free hydrazine is to be determined. Typical results are presented in Table IV. ACKNOWLEDGMENT
The authors wish to acknodedge the assistance of Paul Mohr,
E. A. Brown, and H. A. Gaarder in checking independently the analytical procedures described in this paper. Acknowledgment is also made to the Western Cartridge Company Division of Olin Industries, Inc., for a grant of funds to the University of Illinois to provide the services of part-time analysts in facilitating this investigation. LlTERATURE CITED (1) Bray and Cuy, J . Am. Chem. SOC., 46, 858 (1924). (2) Brown. E. A, thesis, University of Illinois. 1947.
1061
Browne and Shetterly, J . Am. Chem. SOC.,31, 783 (1909). Cuy and Bray, Ibid., 46, 1786 (1924). Dernbach and Mehlig, IND.ENG.CHEM.,ANAL.ED., 14, 58 (1942).
Gilbert, J . Am. Chem. SOC.,46, 2650 (1924). Ibid., 51, 2744 (1929). Ibid.., 58. ~~,1605 ... (1936). Kolthoff, Ibid., 46, 2009 (1924). Kurtenacker and Wagner, Z . anorg. Chem., 120, 261 (1922). Schwarzenbach, H e h . Chim. Acta, 19, 178 (1936). Singh and Rehmann, J . I n d i a n Chem. Soe., 17, 169 (1940). Smith and Wilcox, IXD.ENG.CHEW,AN.LL.ED., 14, 49 (1942). Stolle, J . prakt. Chem., (2) 42, 525 (1890) ; (2) 66, 332 (1902). Saebelledy and Madis, Ber. ungar. pharm. Ges., 13, 368 (1937). ~~
>
~
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.
RECEIVED February 2, 1048. Abstracted from a portion of a thesis submitted by R. A. Penneman t o the Graduate College of the University of Illinois in partial fulfillment of the requirements for the Ph.D. degree.
ANALYSIS OF ALIPHATIC PER ACIDS FRASIC P. GREENSPAN LTD DONiLD G. RIAcKELLAR Riiffalo Electro-Chemical Co., Znc., Buffulo 7, S. Y . A new method for the analysis of per acid solutions is based upon the use of ceric sulfate as a titrant for the hydrogen peroxide present, followed by an iodometric determination of the per acid present. Data are presented showing the results of analyses of sample per acids by the proposed method compared to the potassium permanganate-thiosulfate method of D'Ans and Frey. The new method gires higher, more reliable values.
C
OMMERCI-kL introduction of peracetic acid has stimulated renewed interest in the per acids as bleaching agents, polymerization catalysts, and oxidants in organic synthesis (4, 5, 7 , 10) In the course of extensive work on peracetic acid and other aliphatic per acids, the need arose for an accurate and rapid method of determining aliphatic per acids in the presence of hydrogen peroxide. Hydrogen peroxide is usually found associated with the per acids in aqueous solutions as a result of (1) preparation of the per acid from hydrogen peroxide (8) and (2) hydrolysis of the per acid. Where per acids are prepared by the reaction of an acyl anhydride with hydrogen peroxide, particularly in the absence of a mineral acid catalyst, diacyl peroxide will be present Preparation of per acids by interaction of concentrated hydrogen peroxide and the aliphatic acid, as used by the authors, does not give rise to diacyl peroxides. In a simple and convenient method of analysis of per acid mixtures used by D'Ans and Frey ( 2 , 5) hydrogen peroxide present is titrated in the cold with a standard solution of potassium permanganate to the conventional pink end point. Potassium iodide is then rapidly added and the iodine liberated by the per acid is titrated with a standard thiosulfate solution. This procedure Ras based upon a method of Baeyer and Villiger (1) for hydrogen peroxide-persulfuric acids analyses. Extensive use of the D'Ans and Frey method by this laboratory indicated several inherent deficiencies. The last portion of the permanganate titration was comparativelv sluggish, and this interfered with a sharp and reproducible end point. Variation in time taken before addition of the potassium iodide solution resulted in varying thiosulfate titers for per acid content, and reflected on the accuracy of the method. Furthermore, failure to add the potassium iodide solution immediately a t the permanganate end point gave a rapid development of a deep red-purple color, presumably resulting from the oxidation of Mn-* to a higher valenced manganese compound. This striking and intense color change has been developed by this laboratory into a qualitative test for peracetic acid. D'Ans and Frey ( 2 , originally noted the
oxidation of N n + " by per acid a t the end point and hydrolysis of the per acid as possible sources of error but concluded that such errors were insignificant if one titrated rapidly in the cold. A new analytical procedure for the analysis of aliphatic per acid solutions has been developed, based upon the use of standard ceric sulfate for the initial hydrogen peroxide titration followed by an iodometric determination of the active oxygen present as per acid. Ceric sulfate has been found to satisfy the requirements of a stoichiometrical reaction with the hydrogen peroxide in the presence of the per acid and nonreactivity with the per acid. Further, the C e + + formed does not react with the per acid as does M n + + associated with the use of potassium permanganate. Ceric sulfate has been previously recommended for hydrogen peroxide analyses by Furman and Wallace (6). Diacyl peroxides, if present, do not interfere with the respective hydrogen peroxide and per acid titrations. Such peroxides react very slowly with cold aqueous hydrogen iodide. Where interest lies in a determination of the diacyl peroxide content, this may be obtained by heating the solution being analyzed for 10 minutes on a steam bath after the completion of the hydrogen peroxide and per acid titrations (9). EXPERIMENTAL
Reagents. Ceric sulfate (ammonium tetrasulfatocerate) 0.1 A' in 0.05 h' sulfuric acid. Potassium iodide solution, 10%. Sodium '. Ferroin indicator (0-phenanthroline-ferrous thiosulfate, 0.1 A ion). Procedure. The sample of per acid is accurately weighed and placed in a 500-ml. Erlenmeyer flask containing 150 ml. of 570 sulfuric acid and sufficient cracked ice to maintain a temperature of 0 ' to 10' C. An adequate sample is chosen, when possible, to give approximately 40 ml. of thiosulfate titration. Three drops of ferroin indicator are added and the flask contents are titrated with 0.1 A- ceric sulfate to the disappearance of the salmon color of the indicator. Ten milliliters of the 10% potassium iodide solution are then added and the liberated iodine is tirated with 0.1 N sodium thiosulfate. Starch indicator is added near the end point for the thiosulfate titration. Calculations. = ml. of ceric sulfate X N X 17 2 2 10 X sample weight
ANALYTICAL CHEMISTRY
1062 Table I.
Potassium PermangansteThiosulfate Method Total active H202, Per acid, oxygen,
% 6.86 6.90 6.80 6.78 Mean 6.84 0.05 Mean dev. Performic acid 14.30 Perpropionic 0.60 Material Commercial peracetic acid
Analyses of Per Acid Solutions
% 41.30 41.30 41.22 41.24 41.27 0.04 52.63 6.19
%
11.92 11.94 11.87 11.84 11.89 0.04 13.51 0.82
Ceric SulfateThiosulfate Method Total active HzOz, Per acid, oxygen,
% 6.77 6.78 6.75 6.77 6.77 0.01 14.42 0.62
%
42.10 42.03 42.10 42.15 42.10 0.03 55.91 6.36
% 12.04 12.03 12.04 12.04 12.04 0.00 14.39 0.85
amount of potassium permanganate or ceric sulfate for stoichiometrical reaction with the hydrogen peroxide was added immediately. The results of these experiments are shown in Figure I, where At represents the time interval in minutes between the start of
the hydrogen peroxide titration and t,he potassium iodide addition. Per acid values obtained by the potassium permanganatethiosulfate method decrease rapidly with increasing A t . Values obtained by the ceric sulfate-thiosulfate procedure are constant with a variation in At.
acid
ml. of sodium thiosulfate X N X weight of per acid yoper acid = equivalent 10 x sample weight Iml. of ceric sulfate X N ) 4(ml. of thiosulfate X N ) X 8 Sototal active oxygen = 10 X sample weight Y = normality of reagent RESULTS
Analysis of Per Acid Solutions. Commercial peracetic acid solutions were analyzed by the proposed ceric sulfate-t’hiosulfate method. Results are compared in Table I with those obtained using the potassium permanganate-thiosulfate method procedure of D’Ans and Frey. Per acid values found by the two methods are in disagreement. Careful and repeated analyses indicated differences in per acid values in the order of 20 parts per thousand. In all cases the ceric sulfate-thiosulfate method gave the higher value for per acid. Hydrogen peroxide values are also in poor agreement. The per cent total active oxygen differs significantly, indicating that the errors are only slightly compensating and are intrinsic in nature. Analyses of performic and perpropionic acids (Table I) similarly gave higher values by the ceric sulfate-thiosulfate method with differences of the order of 20 to 55 parts per thousand. Analysis of Hydrogen Peroxide. The discrepancies noted in Table I, extending even to the hydrogen peroxide analysis, indicated the advisability of checking the analysis of hydrogen peroxide alone by the ceric sulfate reagent. Results compared in Table I1 with those obtained by the conventional potassium permanganate method show excellent agreement. EFFECT OF TITRATION TIME ON PER ACID VALUES
In the absence of a pure per acid sample that could be used to evaluate the accuracy of the differing results obtained by the ceric sulfate-thiosulfate and potassium permanganate-thiosulfate methods, recourse was made to a study of the influence of titration time on the per acid values. It had been repeatedly noted, in the use of the potassium permanganate-thiosulfate method, that variation in the time taken before the addition of the potassium iodide solution resulted in varying and nonreproducible thiosulfate titers. This was thought to be directly related to the lower per acid values obtained by this method and to the pronounced oxidation of h h + + by the per acid at the end of the hydrogen peroxide titration. In order to check these points, experiments were conducted with both the potassium permanganate-thiosulfate and ceric sulfate-thiosulfate methods, in which the time between the start of the hydrogen peroxide titration and potassium iodide addition was varied. A large batch of dilute peracetic acid was prepared. A 20-ml. ali uot was accurately introduced into ice cold 5y0 sulfuric acid a n 3 titrated by the usual procedure. The time between the start of the initial titration (hydrogen peroxide titration) and the addition of potassium iodide was accurately controlled. Ice was maintained in the flask throughout the entire experiment. Where At was less than 25 seconds, the premeasured, required
i
I
21
L
2
O
4
6
8
IO
12
14
16
At MINUTES
Figure 1. I.
11.
Effect of Titration Time on Per Acid Values Ceric sulfate-thiosulfate method Potamsium permanganate-thiosulfate method DISCUSSION
.In analysis of results obtained with the potassium permanganate-thiosulfate us. ceric sulfate-thiosulfate method shows a variable and decreasing per acid value obtained with the former, dependent upon the time consumed in the initial titration (hydrogen peroxide titration). In contrast, per acid values obtained with the ceric sulfate-thiosulfate method are independent of titration time over the interval studied. A plot of At against per acid values for the two procedures (Figure 1)shows a coincidence of the two curves a t zero time. Accordingly, it is believed that the zero time value is the true per acid value. The deviations from “zero time” values obtained over a variable titration time for the potassium permanganate-thiosulfate method are shown in Table I11 as derived from Figure 1. Under normal titration procedure, 25 to 45 seconds will be required for the hydrogen peroxide titration. Therefore, errors of 10 to 25 parts per thousand in per acid values are to be expected during the course of a determination by the potassium permanganatethiosulfate method. This agrees with the observation made in Table I that determinations of peracetic acid by the two methods differ by 20 parts per thousand in per acid values. Over the same time interval, no deviations in per acid value are noted with the ceric sulfatethiosulfate method. This method is therefore practical for accurate analysis. Table 11.
-4nalyses of Hydrogen Peroxide Solutions
Hydrogen Peroxide Solution 1
2
Ceric Sulfate Method, 7,
Potassium Permanganate Method,
?& 26.58 26.59 26.57 9.94 9.93
26.55 26.55 26.55 9.95 9.96
Table 111. Effect of Titration Time on Per Acid Assays by Potassium Permanganate-Thiosulfate Method Titration Time, Seconds
Error, Parts per Thousand
20 30
4.5 11 17 25
40
50
V O L U M E 20, N O . 11, N O V E M B E R 1 9 4 8 LITERATURE CITED
(1) Baeyer, A., and Villiger, V., Ber., 34,854 (1901). D'Ans, J., and Frey, W., Ibid., 45, 1845 (1912). D'Ans, J., and Frey, W., Z.anorg. Chem., 84,145 (1913). D'Ans, J., and Kneip, A.,Ber., 48,1136 (1915). Findley, T., Swern, D., and Scanlan, J., J . Am. Chem. sot., 67,
(2) (3) (4) (5)
412 (1945). (6) Furman, N. H., and Wallace, J. H., Ibid.,51, 1449 (1929).
1063 (7) Greenspan, F. P., Ind. Eng. Chem., 39,847 (1947). (8) Greenspan, F. P., J . Am. Chem. Soc., 68, 907 (1946).
(9) Swern, D., private communication. (10) Swern, D., Billen, G., Findley, T., and Scanlan, J., J. Am. Chem. Soc., 67, 1786 (1945). RECEIVEDMay 25, 1948. Presented before the Division of Analytical and Micro Chemistry a t the 113th meeting of the AMERICAN CHEMICAL SOCIETY, Chicago, Ill.
Determination of Small Amounts of Benzene in Ethyl Alcohol D. L. KOLBA, L. R. KANGAS, AND W. W. BECKER Hercules Experiment Station, Hercules Powder Company, Wilmington 99, Del. Small amounts of benzene (0.05 to 1.0% by volume) in ethyl alcohol may be determined by taking advantage of the difference in refractive indices of benzene (1.5014) and hexane (1.3754). The sample of alcohol is diluted with water and partially distilled; the alcohol fraction containing the benzene azeotrope is collected and extracted with hexane. The refractiveindex of the hexane-benzene mixture is measured, and its benzene concentration read from a graph. The difference in the specific gravities of benzene (0.879) and hexane (0.660) provides a second method of measuring the benzene in the hexane extract.
I
K T H E manufacture of solvent types of smokeless powder a t the ordnance plants during World War 11, large quantities of 2B ethyl alcohol, which contains 0.5y0benzene by volume, were used. After its recovery in plant stills, the alcohol contained varying amounts of benzene (0.05 to l.0701, which had to be accurately determined. Some lots rontained up to 1% ether. The methods of D o h ( 4 ) and Baernstein (3) could not be applied directly to the determination of benzene in alcohol. The generally accepted method has been that of Babington and Tingle ( 2 ) . However, when their method was applied to known alcohol-benzene mixtures, high results were obtained, the error varying with the amount of benzene present. It was found that use could be made of the dilution-distillation step of their procedure, in which the benzene comes over in the azeotropic mixture boiling a t 64.9" C. (1). The distillate was extracted with hexane, the refractive index of the hexane-benzene mixture was measured, and the benzene content was then read from a graph prepared by using known mixtures. I t was found also that the measurement of the specific gravity of the hexane-benzene mixture could be used to determine the concentration of the latter. BABINGTON AND TINGLE METHOD
In the method of Babington and Tingle ( 2 ) a diluted sample of the alcohol is distilled, and the first 20-ml. fraction is collected and treated with potassium dichromate This reaction mixture is then extracted with petroleum ether, and the increase in volume of the latter is a measure of the benzene present. This method was applied to several known alcohol-benzene mixtures. The results in Table I show not only that this procedure gives high results, but that the magnitude of the error varies with the amount of benzene present-for example, when the method was applied to pure 190-proof grain alcohol, an apparent benzene content of 0.20 to 0.25% was found. No significant improvement resulted when the method was modified by measuring the decrease in volume of the aqueous layer after extraction. If the chemistry of the method is considered critically, it becomes evident that only a little of the alcohol is oxidized by the dichromate. Only 0.45 gram of potassium dichromate is added, whereas 55.5 grams would be required to oxidize 20
ml. of 80% alcohol to acetic acid. It would seem that the additiog of less than 1% of the theoretical amount of potassium dichromate serves no useful purpose. In several experiments the potassium dichromate oxidation step was omitted. The alcohol was extracted directly with petroleum ether and the latter was extracted with water. The use of a higher boiling hydrocarbon, hexane, was also tested. While fair results were obtained on 2B ethyl alcohol alone, they were too high when the alcohol contained added ether. The measurement of the decrease or increase in volume, therefore, did not appear feasible. REFRACTIVE INDEX METHOD
In the evaluation of the refractive index procedure, known hexane-benzene solutions were prepared and their refractive indices were measured. The plotted values were found t o lie on a straight line. However, when known amounts of benzene were added to 20-ml. pxtion3 of 83% (by w2ight) alcohol and these mixtures were extracted with 10-ml. portions of hexane, the refractive indices of the latter were lower than the calculated values. This was presumably due to dissolved alcohol, which has a refractive index of 1.3624. Accordingly, 25-ml. portions of saturated sodium chloride solution were added to 20-ml. portions of 80% alcohol containing benzene, and these mixtures were extracted with 10-ml. portions of hexane. The refractive indices of the latter agreed closely with the calculated values. A typical refractive index-benzene concentration graph is shown in Figure 1. The alcohol recovered from some plant operations may contain
Table I. Determination of Benzene in Known AlcoholBenzene Solutions by Babington and Tingle Method Benzene Added t o !OO Ml. of Grain Alcohol MI. 0.00 0.20
0.50 0.80 1.oo
Benzene Found Increase in volume of Decrease in volume petroleum ether layer of aqueous layer MI. MI. 0.20 0.40 0.65 0.92 1.15
0.25
...
0.68 0.95
...
