erence species (whose diffusion coefficient must be known). Values in a range of 1.2 < q < 1.4 were obtained for various MFE’s used in this investigation ( 2 ) . This compares with a theoretical assignment of q = d s f o r a model of close packed spheres. The agreement is quite satisfactory since the frits are fabricated by fusing together small particles of glass via a sintering process. Unknown diffusion coefficients can be evaluated at a specified MFE on the basis of Equation 11 through the straight-forward proportionality relationship
where the subscripts X and R identify the unknown and reference species, respectively. The possibility is being explored to utilize the reactions
O2+ 4e
+ 2H20 = 40HO2 + 4e + 4Hf = 2H20
(13a) (13b)
in this context. Reliance on oxygen as the reference species has the following
advantage: ( i , t 1 / 2 ) R , ~can 2 be determined in an air saturated solution, which can subsequently be deaerated with an inert gas. Residual oxygen trapped in the frit can conveniently be removed-and converted to water (or OH-)-by electrolysis in situ. The electroreactive species whose diffusion coefficient is sought can then be added to the same electrolysis cell. I t is anticipated that the accurate determination of polarographic diffusion coefficients can be completed in a few hours in this manner. Since the scarcity of information of this type is due to the intractability (and dubious accuracy) of methods heretofore available ( 4 ) , it is expected that a wealth of information on diffusion coefficients in media of electroanalytical and general electrochemical significance will become available through measurements at the hlFE. Detailed procedures will be described in a full length article at a later date, when it is also planned to discuss interesting applications of the MFE to determination of rate parameters of irreversible electrode processes as well as to investigation of adsorption phenomena.
LITERATURE CITED
(1) Bowers, R. C., Wilson, A. M., J. Am. Chem. SOC.80,2968 (1958).
(2) Clausen,
T . H., Ph.D. thesis, The Pennsylvania State University, 1964;
Dissertalzon Abstr. 26 (2), 660 (1965). (3) Jolley, E. L., Corning Glass Works,
Corning, N. Y., private communication,
November 1963. (4) Macero, D. J., Rulfs, C. L., J. Am. Chem. SOC.81, 2942 (1959). (5) Marple, T . L., Rogers, L. B., A N ~ L . CHEM.25, 1351 (1953). (6) Nicholson, R. S., Shain, I., Ibad., 36, 706 (1964). JUDITH H. CLAWSEN Instrumentation Laboratory 9 Galen St. Watertown, Mass. GERALD B. Moss JOSEPH JORDAN Department of Chemistry 212 Whitmore Laboratory The Pennsylvania State University University Park, Pa. 16802 Received for review June 13, 1966. Accepted June 28, 1966. Based on Ph.D. theses by Judith H. Clausen and Gerald B. Moss. Reported preliminarily before the XIXth International Congress of Pure and Applied Chemistry, London, England, 1963. Supported in part by the United States Atomic Energy Commission under Contract AT(30-1)-2133 with The Pennsylvania State University, Report Number NYO-2133-39.
Electrogenerated Hypobromite SIR: A number of coulometric determinations of hydrogen peroxide have been described. Coulometrically generated Mn (111) (5) has been used to titrate 0.7 to 5.5 mg. of hydrogen peroxide with an average error of 0.3 to 0.4%. The end point was a change of color with ferroin or nitroferroin. (4) employed electroTakahashi generated cerium (IV) in the titration of 50 peq. of hydrogen peroxide with an error of 2%. A more sensitive and precise method was recently developed by Christian (2). He employed coulometric iodometry to titrate 0.05 to 150 peq. of hydrogen peroxide with an average relative error ranging from 0.1 to 4%. In this method the peroxide was added to an iodide solution in the presence of a catalyst, after which a known amount of thiosulfate was added. The excess thiosulfate was then determined coulometrically with electrogenerated iodine. This method, since it is an indirect procedure, removes one of the main advantages of coulometry1400
ANALYTICAL CHEMISTRY
Le., that no standard solution need be prepared. The present paper describes a direct coulometric determination of hydrogen peroxide with electrogenerated hypobromite, which is both sensitive and precise. The end point is determined amperometrically with two polarizable electrodes. The optimum pH for the titration was investigated. As little as 1 keq. of peroxide was determined with a 0.27, error. Arcand and Swift have described the coulometric generation of hypobromite (1). A direct amperometric end point in the coulometric generation of hypobromite was described by Christian, Knoblock, and Purdy (3). EXPERIMENTAL
Reagent grade chemicals were used without further purification. All solutions were prepared using freshly distilled and deionized water. Hydrogen peroxide solutions were prepared by dilution from the 370
strength and were standardized by titrating 20-ml. aliquots volumetrically with standard potassium permanganate using calibrated volumetric pipets. The generating solution consisted of -1M sodium bromide, in a borate buffer. The pH of the borate solution was adjusted with perchloric acid. The generating anode was a 7.5 sq. cm. platinum foil. The cathode was a 1 sq. cm. platinum foil which was isolated from the solution by a glass tube with a fine sintered frit end. Pretreatment of the electrodes consisted of soaking in a 1: 1 solution of nitric acid. The buffer solution served as the catholyte. The indicating electrodes were two platinum foils (1.5 sq. cm.) with 150 mv. impressed between them. The impressed potential was supplied by a Sargent Model XV polarograph. A Sargent current source (Model IV) was used for the hi$her generating currents and a ChrisFeld hlicrocoulometric Quantalyzer (Model 6) was used for the lower currents. The accuracy of the generated currents was checked by measuring the potential developed across a 0.0027, precision resistor with a
again. The excess hypobromite from the previous titration was added to the value obtained. All successive titrations were performed in a similar manner. RESULTS AND DISCUSSION
The oxidation of hydrogen peroxide to oxygen and water serves as the basis for this determination. The excess hypobromite at the end point causes B rise in the amperometric current. A relatively large current developed, because of depolarization of the indicator electrodes, when the peroxide was added to the generating solution. This current decreased as the peroxide was being consumed by the hypobromite. In all the titrations, a minimum was observed before the final current rise (Figure 1). Advantage was taken of this minimum to increase the accuracy of the end point measurement. The sensitivity of the polarograph was adjusted to expand the minimum to full chart width; therefore, the majority of the titration was off scale. However, when the end point was approached, the current rapidly decreased to the minimum after which the abrupt rise in current produced a straight line which decreased the graphical error (Figure 2). The effect of the pH on the determination was studied. At a pH of less than 7.8, less hypobromite was formed; therefore, a corresponding increase in the titration time was shown. The rounded titration curves a t pH 9 and above, reduced the accuracy of the titration. A narrow range of pH was studied (Table I) in more detail to determine the optimum value. The best results were obtained a t a pH of 8.18-8.30. Since the pH increases slightly during each titration, a value of 8.20 was chosen as the optimum initial pH.
30 SEC.
H
Seconds at 9.65ma. Figure 1. Coulometric titration of 14 peq. of hydrogen peroxide Sensitivity of polarograph, 0.1 Ma./mrn.
Leeds & Northrup volt potentiometer. All titrations were performed in a 50- or 100-ml. beaker using the polarograph to record indicating current with the chart paper calibrated in terms of seconds. The solution was stirred with a magnetic stirring bar during generstion and current measurement. Procedure. A pretitration procedure was used in all analyses. A small amount of peroxide (?50% of sample size) was added to the cell and was titrated to a slight excess of hypobromite. The excess of hypobromite after the end point, was noted. Then the sample was added and titrated until the current rose
Table I.
Effect of pH on Coulometric Determination of Hydrogen Peroxide
No. of
PH
Current
titrations
Added
8.00 8.13 8.18 8.21 8.30 8.41
9.650 4.825 4.825 4.825 4.825 4.825
4 5 5 4 4 4
9.098 9.098 9.098 9.098 9.098 9.098
Table II.
9
Found 9.17 9.140 9.109 9.105 9.090 9.120
Mean error,
Std. dev.
0.79 0.46 0.12 0.076 0.087 0.24
0.064 0.049 0.053 0.021 0.051 0.059
I-
z w
a a 3 u
30sec.
H
SECONDS AT 9.65 ma. Figure 2. Coulometric titration of 14 Meq. of hydrogen peroxide Sensitivity of polarogroph, 0.004 Ma./mm.
The accuracy and precision of the method were determined by titrating hydrogen peroxide which had been previously standardized with potassium permanganate. The recoveries of known amounts of hydrogen peroxide titrated a t different currents are summarized in Table 11. This determination can be applied to a wide range of hydrogen peroxide concentrations. Amounts from 1.O to over 100 peq. of peroxide have been titrated with errors less than 0.20a/0.
( 1 ) Arcand, G. M., Swift, E. H., ANAL. CHEM.28, 440 (1956).
(2) Christian, G. D.. Zbid.. 37. 1418
FREDRIC J. FELDMAN ROBERT E. BOSSHART
Mean error,
peq.
0.08~0.
LITERATURE CITED
%
Recoveries of Hydrogen Peroxide
No. of PH 8.21 8.21 8.21
Meq
I
Current
titrations
Added
Found
%
Std. dev.
4.825 9.650 19.30
14 13 14
7.095 14.48 21.68
7.109 14.46 21.72
-I-0.20 -0.14 +o. 18
0.017 0.022 0.032
Division of Biochemistry Walter Reed Army Institute of Research Walter Reed Army Medical Center Washington, D. C. 20012 USE of manufacturer’s name does not constitute an official endorsement by the U. S. Army. VOL 38, NO. 10, SEPTEMBER 1966
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