Table 1.
Determination of Manganese in Blended Gasolines
Type of Manganese, Grams/Gallon Gasoline Added Found Difference Olefinic 0.052 0.051 -0.001 0.101 0.464 0.757 1.01 2.02 Aromatic 0.066 0.102 0.454 0.755 1.04 2.02
+0.002 +0.001
0.103 0.465 0.758 0.998 2.03 0.064 0.100 0.455 0.755 1.03 2.05
+0.001 -0.01 +O.Ol -0.002 -0.002 +O.OOl 0.000 -0.01
+0.03
Table It. Determination of Manganese in Commerciai Gasolines Containing (Methyfcyclopentadienyl)manganese Tricarbonyl
Gasoline A
Manganese, Grams/Gallon Added Found 0.518
B
0.501
C
0.533
D
0.508
E
0.242
F
0.249
G
0.225
H
0,235
I
0.282
J
0.247
0.520 0,523 0.508 0.512 0.528 0.530 0.508 0.512 0,239 0.242 0,249 0.249 0.223 0,224 0.234 0.234 0.275 0.278 0.247 0.247
nor by the hydrochloric acid extraction used for tetraethyllead (1). Decomposition of AK-33X with bromine, as
used in the determination of trace lead (d), was found effective, A solution of bromine in carbon tetrachloride was the best reagent for decomposing AK-33X in gasoline. This reagent can be used without hazard even in highly olefinic gasolines, if the sample size does not exceed 5 ml., although the recovery of manganese may be incomplete for larger samples of these gasolines. However, the sample limit can be extended to 10 ml. for highly aromatic gasolines, since aromatics are less reactive than olefins. Some of the bromine added to decompose AK-33X reacts with highly olefinic gasolines. To test whether this reaction might be a source of error, a highly olefinic gasoline and a highly aromatic gasoline containing equal amounts of manganese as AK-33X were analyzed by the proposed method (Table I). The olefinic gasoline contained 42% olefins and 14% aromatics. The aromatic gasoline contained 1% olefins and 41% aromatics. These gasolines, blended with AK-33X a t concentrations varying from 0.052 to 2.02 grams of manganese per gallon, gave equally good results. Water can be used to extract the manganese salts from gasoline after bromination. However, acidic solutions are preferred, since they dissolve some of the decomposition products and provide a cleaner separation. Twenty per cent solutions of sulfuric, nitric, and phosphoric acids were equally effective. Each acid removed approximately 99% of the manganese in the first extraction. However, use of sulfuric acid with leaded gasolines produced a white turbidity that interfered with spectrophotometric readings. The phosphoric acid was arbitrarily chosen in preference to the nitric acid. Two determinations on a gasoline sample can be made in 75 minutes; 32 determinations, per man-day. At-
tempts to shorten the analysis time by oxidizing the manganese without first removing the organic material gave poorer accuracy. Moreover, the time required for the removal of organic material when multiple samples are analyzed is negligible, since this oxidation procedure requires little of the analyst’s time. Ten representative commercial gasolines containing known amounts of AK-33X were analyzed in duplicate (Table 11). For manganese concentrations of 0.5 gram per gallon, duplicate results agreed within 0.4 to 0.80/0. For concentrations of 0.25 gram per gallon, duplicates agreed within 0.4 to 1.3y0. Mean values differed from the amount added by 0.0 to 1.8%. An estimate of the precision was obtained by analyzing blends of a commercial fuel containing 0.102, 0.266, 0.464, and 0.755 gram of manganese per gallon. Each blend was analyzed seven times under the same experimental conditions. Standard deviations of 0.0019, 0.0026, 0.0033, and 0.0036 gram of manganese per gallon, respectively, were obtained. LITERATURE CITED
Am. SOL Testine Materials. “ASTM Standards on Petyoleum Products and Lubricants,” Method D 526-56. (2) Griffing, M. E., Rozek, A., Snyder, . L. J., HGnderson, S. R., ANAL.CHEM.
(1) \-I
29, 190 (1957). ‘ (3) Jones, R. A., Ibid., 31,1341 (1959). 1.2) -, Mehlie. J. P.. IND.ENQ. CHEM.. ANAL.I% 11,.274 (1939). (5) Tightingale, E. R., Jr., Wilcox, G. W., Zielinski, A. D., AN&. CREM.32, 625 (1960). (6) Smith, G. W., Palmby, A. K., Zbid., 31,1798 (1959). (7) Willard, H. H., Greathouse, L. H., J . Am. Chem. SOC. 39,2366 (1917). \
RECEIVED for review February 18, 1960. Accepted September 23, 1960.
The Formation of Hydrogen Peroxide in Alcohols Its Effect on the Extraction of Chelates of Benzohydroxamic Acid CLIFTON E. MELOAN’ and WARREN W. BRANDT Department of Chemistry, Purdue University, lafayette, Ind.
b The benzohydroxamic acid chelates of uranium(VI), vanadium(V), and iron (Iff) are destroyed when they are extracted into alcohols which have not been freshly distilled. This work demonstrates that traces of hydrogen peroxide formed in the alcohol are the cause of this effect. A colorimetric 102
ANALYTICAL CHEMISTRY
method using titanium is utilized for the detection and determination of the hydrogen peroxide. The formation of hydrogen peroxide may b e eliminated satisfactorily by saturating the alcohol with water and storing it in a metal container. The formation of hydrogen peroxide was detected in
-
3-methyl-1 -butanol, 1-pentanol, 1 hexanol, 1 -heptanol, 1 -octanol, and cycfohexanol.
ipresent addrees, Kansas State University, Manhattan, Kan.
T
HE BEXZOHYDROXAMIC ACID METBODS for determining uranium(VI),
vanadium(V), and iron(II1) involve the extraction of the chelates into an immiscible alcohol, usually 1-hexanol (8, 10). 1-Hexanol that has not been freshly distilled gives lower absorbance values than that which has been freshly distilled. The amount of this interference increases with time of standing. Some 1-hexanol solutions only 2 to 3 weeks old completely destroyed the uranium(VI), vanadium(V), and iron (111) chelates of benzohydroxamic acid. This effect was also observed to occur with several other saturated alcohols from Cs to Clo. Some compound was being formed in alcohols upon standing that was destroying these chelates and not just preventing the extraction. EXPERIMENTAL
Reagents. Solutions of reagent grade NaV03, UOz(NO&. 6 H z 0 , K2TiO(CzO&. 2H20, and HZOZ (all Fisher Scientific Co.) were used. The benzohydroxamic acid (Aldrich Chemical Co.) was used as a 0.01M aqueous solution. Apparatus. Beckman Model B spectrophotometer equipped with four matched 1.00-em. cells. Leeds & Northrup p H meter (lineoperated). Procedure. The uranium chelate used in these determinations was prepared as follows: Two milliliters of 10+M uranyl nitrate were added to 20 ml. of 10-2M benzohydroxamic acid and diluted to about 50 ml. with distilled water. The pH was adjusted to 6.0 (pH 3 for vanadium chelate). Ten milliliters of the alcohol being tested were used for extraction and absorbance was measured a t 380 mp. The usual determination of titanium with hydrogen peroxide (5) was reversed. This method proved capable of detecting 1 part of peroxide per 100,000. The p H of the titanium solution was adjusted to 1.1 with HC10,. The peroxide was extracted from the contaminated alcohol by mixing equal volumes of the standard titanium solution and the alcohol. A very deep yellow color, due to the titanium peroxide complex, was produced in the aqueous layer, the absorbance of which was measured a t 370 muL1. RESULTS AND DISCUSSION
When the titanium solution was mixed with alcohols that would not extract the chelate, a very deep yellow color of the titanium peroxide complex was produced proving that peroxides were present in the alcohols. The interference of hydrogen peroxide was confirmed by adding small amounts of hydrogen peroxide to the benzohydroxamic acid chelates of uranium(VI), vanadium(V), and iron(II1). In all cases the chelate was destroyed im-
0
IO
20
30
40
Days
Figure 1. Chelate destruction by and peroxide formation in 1 -hexanol
mediately. Other qualitative tests indicated the absence of significant amounts of aldehydes, hydroperoxides, and alkylperoxides as causes of the interference. Specificity of Titanium for Hydrogen Peroxide. The yellow color formed by the titanium reaction is not due to compounds of the type ROOH and ROOR (9). Babko and Volkova (1) have shown that the colored species of the titanium and hydrogen peroxide is a 1 to 1 complex. A calibration curve of the reversed system follows Beer's law over the concentration range discussed. The only other type of peroxide compound with which titanium will form a yellow complex is an hydroxyl alkyl peroxide, RCN(0H)OOH (9). Certain inorganic peroxides will also produce a yellow color, but only if they can be converted to hydrogen peroxide during the course of the analytical procedure (5, 7). The organic peroxides are of more concern and the mechanism for their formation reported by Sutterfield and Bonnel is: HCHO HCHO
+ Hz02 HCH(0H)OOH + HCH(0H)OOH e i~
(1)
HCH(OH)OOCH(OH)H (2) The equilibrium in both equations is shifted to the left if a large excess of acid is present. The first equation is important because the product gives a yellow-colored titanium complex while the product of the second equation does not. Infrared examination showed that aldehydes are present in the rontaminated alcohols. The reversal of Equations 1 and 2 in acidic solution requires about 2 hours a t room temperature, if the pH is lower than 2. As a check on the formation of HOOH from RCH(OH)OOH, ROOH, and ROOR, two series of titanium complexes were prepared using contaminated 1-hexanol. The absorbance of one series was measured immediately and the other was measured 2l/2 hours later. There was an increase in intensity in the second set over the first which indicates that about
5% total-RCH(OH)OOH, ROOH, or ROOR may have been present. This amount is not considered significant in this study. Peroxide Formation in Alcohols. Bolland and Cooper (2) found hydrogen peroxide among the products of a photochemical oxidation of alcohols using a sensitizer. Only 0.002 mole (0.068 gram) of hydrogen peroxide in 1 liter of 1-hexanol would completely destroy the chelate of benzohydroxamic acid. Such a small amount may have been overlooked previously. It is difficult to postulate a free radical mechanism for producing compounds of the type ROOR and ROOH without also producing HOOH. Peroxide-Chelate Relationship. A stoichiometric relationship was established between the peroxide formed and the chelate of benzohydroxamic acid destroyed by extracting the uranium(V1)-benzohydroxamic acid chelate with 1-hexanol containing varying amounts of peroxide (Figure 1). The decrease in absorbance of the chelate was compared with the amount of color produced by standard solutions of hydrogen peroxide. A plot of the increase in absorbance of the titanium peroxide vs. the decrease in absorbance of the chelate produced a straight line terminating with a break a t the point where 1.92 x 10-6mole of the hydrogen peroxide completely destroyed 2.00 X 10-6 mole of the uranium(V1)--benzohydroxamic acid. This shows within experimental error a 1 to 1 ratio between the chelate destroyed and the peroxide formed. These results are in agreement with the work of Corpel (4), who has found a 1to 1 ratio between the uranyl ion and hydrogen peroxide existing up to a p H of 10.3. Prevention of Peroxide Formation. The alcohols were stored over iodine, activated alumina, and various oxidizing, reducing, and drying agents u i t h little inhibition of peroxide formation being noticed. Ultraviolet radiation was very important to the formation of peroxides and for this reason a metal container is recommended for the storage of alcohols. Darkened bottles were not satisfactory. The effect of water on the formation of peroxide was determined. 1-Hexanol was distilled to purify it and then solutions containing up to 7% water (saturated) were prepared. They were placed in clear glass bottles and stored in the sunlight. Their degree of contamination was periodically tested by the previously described methods of uranium chelate destruction and titanium peroxide formation. The results are shown in Figure 2. The peroxide is either destroyed or does not form in the presence of water. Based upon these studies the most satisfactory way to keep a peroxide-free VOL 33, NO. 1, JANUARY 1961
103
r-7?7
r " " ;added during the distillation, the
Days
Figure 2. Chelate destruction (upper) and peroxide formation (lower) by 1-hexanol containing different amounts of water
alcohol for the purposes of extraction is to store it in a metal container and to saturate it with water. Removal of Peroxides Formed. The alcohols purchased commercially are usually contaminated with traces of peroxides unless they have been stored in metal containers. A distillation of the contaminated alcohol will remove the peroxides; however, if anhydrous potassium carbonate is
absorbance values obtained in the subsequent extraction are about 5% higher, indicating that the anhydrous potassium carbonate has removed more of the hydrogen peroxide. The purpose of the potassium carbonate could be to provide potassium hydroxide from the traces of water present. Potassium hydroxide will destroy hydrogen peroxide (6). Therefore, distillation of the 1-hexanol over potassium hydroxide should be much more effective than just potassium carbonate (Figure 3 shows comparative results). Using the undistilled alcohol as a reference solution which forms peroxides a t a certain rate (curve D, Figure 3), the distilled alcohol which has had most of the hydrogen peroxide removed should have less peroxide initially (curves A , B, C). However, the distilled l-hexanol should have less water present and therefore produce peroxide faster. This is also shown to be the case. These effects account for the lower initial value and the succeeding faster formation rate of the hydrogen peroxide. ACKNOWLEDGMENT
The authors express their appreciation for the financial support of this work by the Purdue Research Foundation. LITERATURE CITED
(1) Babko, A. K., Volkova, A. I., J . Gen. Chem. U.S.S.R. (Eng. Transl.) 21, 2163 /14R1) \-.,.,-,. (2) Bolland, J. L., Cooper, H. R., Nature 172,413 (1953). (3) Bonet-Maury, P., Compt. rend. 218, 117 (1944). (4) Corpel, J., Bull. SOC. chim. France 1953, 752.
I6-
I
I
I
I
I
0
3
6
9
12
I
Hours
Figure 3. Peroxide formation in 1 -hexanol irradiated with a mercury vapor lamp A. 5. C.
D.
After distillation over KOH After distillation After distillation over KZCOB Not distilled
( 5 ) Furman. 15'1 Furman, N. H.. H., "Scott's Standard \ - I
Methods 'of of Chemical Analysis," 5th ed., Vol. I, p. 987, Van Nostrand, New York, 1947. (6) Kasosnoskii, I. A,, Neidig, A. B., Dokladu A k a d . Nauk S.S.S.R. 86. 717 (1952)." (7) Klenk, J. v., Klepzig's Teztil. 2. 42, 549 (1939). (8) Meloan, C. E., Brandt, W. W., Holkeboer, P. E., ANAL. CHEM.32, 701 /iacn\ (9) Sutterfield. C. N.. Bonnel. A. H.. Ibid., 27,1174 (1955): (10) Wise, W. M., Brandt, W. W., Ibid., 27,1392 (1955). "I \L""",.
RECEIVED for review May 17, 1960. Accepted October 17, 1960. Abstracted from a thesis submitted by Clifton E. Meloan t o the Graduate School of Purdue University, January, 1959, in partial fulfillment of the requirements for the degree of doctor of philosophy.
Spectrophotometric Determination of Traces of Hydrogen Peroxide C. E. MELOAN, M. MAUCK, and C. HUFFMAN Department of Chemistry, Kansas State University, Manhattan, Kan. Hydrogen peroxide reacts with the colored benzohydroxamic acid chelates of vanadium(V) and uranium(V1) to destroy the chelate and form the metal peroxide complex. The decrease in color of the benzohydroxamic acid chelate is a measure of the amount of peroxide present. It i s possible to mole of hydrogen detect 7 X peroxide using the uranium chelate (pH 6 and 380 mp) and 1 X 10+ mole using the vanadium system (pH 3 and 450 mp). By using a 1 hexanol extraction the color due to
104
ANALYTICAL CHEMISTRY
the metal peroxide complexes formed does not interfere. Hydroperoxides and dialkyl peroxides do not interfere appreciably at the pH's required for this determination.
W
HEN
EITHER
THE CHELATES Of
vanadium and uranium with benzohydroxamic acid (6, 9) are extracted with 1-hexanol, a large decrease in absorbance is observed in some cases. Meloan and Brandt (7) have shown this to be due to the presence of traces
of hydrogen peroxide in the alcohol. This chelate destruction is proposed as a sensitive and selective method for the determination and measurement of hydrogen peroxide. Methods currently in use for the determination of traces of HzOz are the reversal of the standard colorimetric procedure for titanium using H ~ O Z (81, the use of luminol (I),and polarography (4). While each of these methods are in general satisfactory, each one has some undesirable characteristics. The tita-
