825
V O L U M E 28, NO. 5, M A Y 1 9 5 6 minations for portion 2 are uncertain because of the small quantities of ozone involved. Within the limits of experimental error, however, the spectroscopic data appear t o confirm that all of the atmospheric oxidant desorbed from silica gel is ozone. ACKNOWLEDGMENT
The work described in this paper was carried out at the Pasadena Laboratories of Stanford Research Institute under the sponsorship of the \T7estern Oil and Gas Association. Specti oscopic analyses were carried out a t the Brea Research Center of' the Union Oil Co. The authors wish to express their thanks t o \T, C. Merrill and Gene Schluter, who performed the spectroRropic analyses.
LITERATURE CITED (1) Air Pollution Foundation, Los Angeles, Calif., Rept. 12 (1955). (2) Bartel, -4. W., Temple, J. W., Ind. Eng. Chem. 44, 857 (1952). (3) Bradley, C. E., Haagen-Smit, 9. J., Rubber C h e n . and Technol. 24, 750 (1951). (4) Crabtree, J., Biggs, B. S., J . Polymer Sci. 1 1 , 280 (1953). (5) Crabtree, J., Kemp, A. R . , I n d . Eng. Chem. 38, 278 (1946). (6) Edgar, J . L.. Paneth, F. A , , J . Chem. Soc. (London) 1941, 511. ( 7 ) Haagen-Smit, A. J . , "Chemical ilnalysis of Air Pollutants," Air
Pollutant Proc., U. S. Technical Conference on Air Pollution, p. 193, AIcGraw-Hill, S e w York, 1950. (8) Littman, F. E., Benoliel, 11. W., A N . ~ LCHEM. . 25, 1480 (1953). (9) Regener, V. C., "Atmosphere Ozone in the Los Angeles Region," Univ. New Mexico, Contract 4 F 19(122)-381, Sci. Rept. 3 (July 22, 1954). R E C E I V Efor D review December 2 , 1955. Accepted March 5 , 1956
Polarographic Study of Alkyl Hydroperoxides D. A. SKOOG and ALLEN B. H. LAUWZECHA Department of Chemistry, Stanford University, Stanford, Calif.
The polarographic behavior of alkj 1 hydroperoxides w a s investigated systematically in order to define the most suitable conditions for their determination. Se\enteen examples of these compounds varying from four to nine carbon atoms in chain length were investigated. 4queous alcohol solutions of various compositions were used as solvents in this work. Several supporting electrolytes were investigated, and, of these, sulfuric acid was found to be the most satisfactory. The best polarographic waves were obtained with the compounds of higher molecular weight. These were steeper, better defined, and occurred at a less negative potential. The half-wave potentials for all of the hydroperoxides were so close together that the determination of anj indi\idual compound in a mixture was impossible. In general, diffusion currents were directly proportional to concentration. However, the proportionalitj constants for the barious hydroperoxides \aried widelj, limiting the usefulness of the method for functional group analysis.
I
N R E C E X T years there has been considerable interest in
methods for the determination of the various types of organic peroxides. The polarographic procedure is among the several methods n-hich have been suggested for this. However, only a relatively small number of examples of organic peroxides have been investigated because of the difficulty in preparation of pure samples of many of these compounds. As a result of the recent work of Williams and hlosher ( I O ) , it is possible to prepare the alkyl hydroperoxides of lox-er molecular weight in a fairly pure state, and it seemed worth while to study the polarographic behavior of a number of these in order to establish conditions best suited for their determination. Several references are found in the literature to the polarographic reduction of the first two members of the alkyl seriesnamely, methyl and ethyl hydroperoxide (2, 6, 7 , 9). I n addition Willits and others (11) have reported on the reduction of 3-butyl hydroperoxide as well as several high molecular weight compounds including cumene, menthane, pinene, and cyclohexane hydroperoxide. Their work was done in a nonaqueous solvent of methanol and benzene Lvith lithium chloride as the supporting electrolyte. RIacXevin and Urone ( 5 ) ,in a paper describing a method for the determination of hydrogen peroxide in the presence of hydroperoxides, show a polarogram for 3-pentyl
hydroperoxide from a 0 . 1 s solution of potassium chloride. Bruschweiler and Minkoff ( I ) have recently repoited on the reduction of methyl, ethyl, and %butyl hydroperoxide from lithium sulfate and lithium hydroyide solutions. This paper describes a systematic study of the reduction of a number of typical alkyl hydroperoxides varying from four t o nine carbon atoms in chain length. Water and aqueous ethyl alcohol solutions vere used as solvents in this work and several supporting electrolytes x ere investigated. APPARATUS AND MATERIALS
-411 of the polarograms reported herein were recorded Ivith a Sargent Model XXI Polarograph. An H-type cell similar t o t h a t described by Lingane and Laitinen ( 4 ) was used throughout this work. The cell was thermostated a t 25' i 0.1" C. The dropping mercury electrode was of the usual type and had a capillary constant, m2'3t1'6,of 2.644. K i t h the exception of 3-butyl hydroperoxide, which was obtained commercially, all of the hydroperoxides were prepared according t o the method of Williams and Mosher ( I O ) , and some of the compounds studied were furnished by them. Table I1 lists the 17 alkyl hydroperoxides which were investigated. Based on iodometric titration data, the purity of all of the compounds listed is believed t o be a t least 95% and in most cases better than 99%. Solutions. Sulfuric Acid, 0.5F. Potassium Hydroxide, 0.5F. Gelatin Solution, 0.1%. Methyl Red, 0.401,. BUFFERSOLUTIOSS.Approximately 330 ml. of 0.2F acetic acid was mixed with 20 ml. of 0.2F sodium acetate; 35 ml. of this was used for each 100 ml. of solution prepared. This solution buffers at approximately pH 4.0. For a buffer a t approximately pH 9.0, about 53 ml. of 0.1F potassium dihydrogen phosphate was mixed with 347 ml. of 0.05F borax; 40 ml. of this was used for each 100 ml. of solution prepared. STOCKSOLL TIOSS OF PEROXIDES.Approximately 0.1F stock solutions of the peroxides were prepared by diluting carefully weighed portions of the compounds to exactly 25 or 50 ml. with 95% ethyl alcohol. These solutions were stored in brown bottles sealed with aluminum foil and kept a t -10" C. except when used. Such solutions were found t o be stable for several months under these conditions. STASDARD PEROXIDE SOLUTIONS. The stock peroxide solutions were allowed t o come t o room temperature and measured volumes were then transferred t o 100-ml. volumetric flasks. The supporting electrolyte and sufficient 95% ethyl alcohol to give the desired concentration of this reagent we?e added and the solution was diluted t o the mark. The alcohol concentrations referred to in this paper are volume percentages of the 95% ethyl alcoholLe., milliliters of 95% ethyl alcohol per 100 ml. of solution.
ANALYTICAL CHEMISTRY
826 60
% E T H Y L ALCOHOL
Table I. Effect of Supporting Electrolyte on Polarograms of Hydroperoxides Hydroperoxide
H?Soa, 0.3F
H a l f - T a r e Potential, Volt us. 8. C . E. Acetate Phosphate buffer, buffer, KOH, pH 9 0.1.\pH 4
1 Butyl
-0.26
-0.21
-0.20
3-Butyl I-Pentyl 2-Pentyl
-0.35 -0.20 -0.20
-0 2 i -0.19 -0.15
-0.27 -0.20 -0.14
Drawn out warebetween - 0 . 3 and - 1 . 3 roira
......
Draivn out wave ,.....
50 0
c a
5
30
Bi: I r
3.0
C
0 3
PROCEDURE
T h e solutions were transferred to the polarographic cell, brought to 25" C., and freed of oxygen by bubbling nitrogen through the cell for 10 to 20 minutes. T h e nitrogen x a s first passed through a tower containing a solution similar to that being studied to minimize composition changes due to evaporation. T h e solutions were blanketed with nitrogen during the run to prevent absorption of oxygen. All of the polarograms were recorded a t a rate of potential increase of 0.118 volt per minute. I n calculating half-wave potentials no correction for iR drop was made. The cell resistance throughout this n-ork was 350 to 450 ohms, and generally such corrections would be smaller than the accuracy x i t h which potential measurements were made. Furthermore, the uncertainties introduced by variations in the liquid junction potential with solvent composition were probablv of the same order of magnitude as the variations in iR drop l\-ith composition. EXPERIMEZITAL RESULTS
Solvent. Of the compounds investigated only the butyl hydroperoxides ivere found to be sufficiently water-solutile for convenient study in this medium. Aqueous solutions of ethj.1 alcohol, on the other hand, readily dissolved all of the compounds under investigation, and mixtures containing from 5 to 60% ethyl alcohol by volume were used. T h e concentration of ethyl alcohol was found to have a profonnd effect on the polarograms for the hydroperoxides, particularly the lower molecular weight members of the series. The effect is illustrated in Figures 1 and 2 . I n general, an increase in the alcohol concentration led to less well-defined. and more dravin out waves with smaller slopes. Half-mve potentials were more negative and diffusion currents smaller a t higher alcohol concentrations. A similar variation in behavior for the reduction of p-nitroaniline in ethyl alcohol-Tvater mixtures has been discussed by Shreve and Markham (8). T h e higher molecular n-eight peroxides gave satisfactory waves from solutions containing 10% or more alcohol. However, to obtain \Tell-defined waves with t,he pentyl peroxides, the alcohol Concentration had t o be kept below 2075, and with the butj-1 peroxides below 10%. For analytical purposes, the alcohol concentration of the solvent should be kept' as low as possible. Supporting Electrolytes. Several types of supporting electrolytes were investigated. These included 0.1S potassium hydroxide, a phosphate buffer of pH approximately 9, an acetate buffer of pH approximately 4, and sulfuric acid varying in con centration from 0.1F t o 0.5F. I n general the most satisfactory waves were obtained from the sulfuric acid solutions, although the polarograms from the phosphate and acetate buffer solutions were nearly as satisfactory in most cases. With 0.LV potassium hydroxide, extremely drawn out and analytically unsatisfactory waves Tvere obtained. In Table I are listed data obtained for four of the hydroperoxides in the various supporting electrolytes. It will be noted that the half-wave potentials become slightly less negative with increasing pH; however, this effect is comparatively small. Variation in the supporting electrolyte had a negligible effect on the diffusion currents obtained.
2.0
10
I
00
-01
-03
-05
-07
-09
-1 1
APPLIED POTENTIAL, VOLTS v s S C E
Figure 1.
Effect of alcohol concentration on polarograms for 1-butyl hydroperoxide
Sol4:ions 5 X 10-rF in hydropero:,ide and 0.5F in sulfuric acid
For purposes of analysis, sulfuric acid appeared to be well suited as a supporting electrolyte. Less trouble with maxima was encountered with this than Tvith the other electrolytes, and satisfactory Tvvaves n-ere obtained for all of the perovides studied. J t was found that the concentration of sulfuric acid could be varied from 0.1 to 0.5F without appreciably affecting the polarograms either as to half-wave potential or diffusion current, and most of the data reported in this paper are for solutions of sulfuric acid in this concentration range. Effect of Molecular Weight and Chain Branching on Polarograms of Hydroperoxides. Table I1 gives half-wave potentials and diffusion current constants, i d / C , for all of the hydroperoxides qtudied. The data obtained mere for solutions 0.1F in sulfuric acid and, with the exception of the butyl compounds, 20% b\ volume in ethyl alcohol. The half-n-ave potentials became less negative with increasing molecular 11-eight. Furthermore, there appeals to be a slight negative shift in half-wave potentials in going from the normal isomers to the secondary and tertiary compounds. However, none of the half-wave potentials is sufficiently different t o allow analysis of any particular hydroperoxide in a mixture of the compounds.
Table 11. Half-Wave Potentials for Various Alkyl Hydroperoxides Hydroperoxide 1-Butyl %Butyl 3-Butyl 1-Pentyl %Pentyl 3-Pentyl 3-Methyl-1-butyl Cyclopentyl I-Hexyl 2-Hexyl 3-Hexyl Cyclohexyl 1-Heptyl 2-Heptyl 1-Octyl 2-Octyl I-Nonyl
Alcohol, Vol. 7 0
?
?,
20 20 20 20 20 20 20 20 20 20 20 20 20 20
Euz, Volt 3 8 . S.C.E. -0.26 -0.28 -0.34 -0.20 -0.24 -0.22 -0.23 -0.25 -0.12 -0.16 -0.16 -0.14 -0.03 -0.12 -0.02 -0.80 -0.01
%d/C
8.4 8.2 8.2 7.8 7.0 6.7 7.0 7.0 6.9 6.7 6.6 6.7 6.3 6.0 6.0 5.8
V O L U M E 28, NO. 5, M A Y 1 9 5 6
827
The diffusion current constants also decrease appreciably with increasing molecular weight. Probably, this decrease simply reflects the decrease in diffusion rate which ould accompany the increase in the size of the molecules. The shapes of the polarograms vary considerably in going from the lower molecular weight compounds to the higher (Figure 3). With the butyl and pentyl hydroperoxides, the waves are drawn out and rather poorly defined, m-hereas the octyl and nonyl compounds yield steep, sharp curves which are in every way more satisfactory. It is apparent that with the smaller molecules, the electrode process is a highly irreversible one. K i t h the octyl and nonyl compounds, on the other hand, the slopes of the waves approach those for a two-electron reversible reduction. Plots of log i / ( i d - i) against potential for these were straight lines with slopes of 0.026 and 0.024, respectively, compared with a theoretical slope of 0.030 for a two-electron reduction. However, the independence of these vaves of p H changes indicates t h a t the electrode reaction is not reversible even with the highest molecular weight compounds studied.
From the data in Table 111 and other similar data, it v-as concluded that it is possible t o determine any one of the hydroperoxides studied with an accuracy of 1 t o 2 % by the polarographic method. However, the rather large variation in the diffusion current constants ( i d / C ) among the various hydroperoxides seriously limits the accuracy of the method for a functional group analysis, a t least for hydroperoxides in the range of 4 to 9 carbon atoms. Thus, in Table I1 i t m-ill be seen that a 20% decrease in i d / C is found in going from the 1-pentyl (n-amyl) to the 1-nonyl compound.
Table 111. Hydroperoxide 3-Butyl
l-Pent,l
I-Octyl
Effect of Concentration on Polarograms
Solvent a n d Hydroperoxide Supporting Concn., C, Electrolyte Mmoles per Liter 0.30 5 % ethyl 0.50 alcohol, 1.00 0 3F 2.00 HzSO4 3.00 sp0 ethyl 0.30 0.50 alcohol, 0 3F 1.00 2.00 H?SO4 3.00 40% ethyl 0.30 alcohol, 0.50 0.3F 1.00 HzSOa 2.00
Diffusion Current, id, Fa.
2.5 4.2 8.1
16.2 24.5 2.5 4.2 8.5 16.5 24.; 1.7
;:;
11.2
___~
Eli?,
Volt us. S.C.E. -0.30 -0.30
id/C 8.3 8.4 8.1 8.1 8.2 8.3 8.4 8.5 8.3 8.2 5.G
-0.32
-0.31 -0.32 -0.18 -0.18 -0.18 -0.18 -0.19 -0 06 -0 07
5.i
-0.06 -0 05
3.5
5.6
-
~
~-
Stability of Solutions. The stock solutions of the hsdioperoxides in 95y0 ethyl alcohol were found to be stable for several months at -10” C. The more dilute aqueous solutions containing the peroxides and suppoi ting electrolyte were stable for at least a n-eek \)-hen refrigerated, but at room temperature and i n contact with mercury, slov- decomposition of these solutions took place. The rate of decomposition, however, was don- enough to cause no serious difficulties in obtaining reproducible polarogranis. For example, with a 0.001F solution of 2-but)-1 hydroperoxide in sulfuric acid, a decrease in current from 7.95 to 7.6 pa. was observed in 1 hour of standing in contact n-ith mercury. GO ‘02
1
-00
-02
-04
- c-
- c
I
BC
-2 5
APPLIED P O T E N T I A L , LOLTS v s S C C 7.c
Figure 2. Effect of alcohol concentration on polarograms for 1-octyl hydroperoxide Solutions 5 X 10-‘F in hydroperoxide and 0 . 5 F in sulfuric acid
Current Maxima. Current niavima mere encountered from time t o time in this n-ork. I n general, these appeared more frequently when TT orking a t lower alcohol concentrations and particularly mith the higher molecular weight peroxides. The alcohol appeared to act as a suppressor and maxima were seldom found when alcohol concentrations of 20% or greater were used. Maxima occurred more frequently and vere larger in the acetate buffer than in either the sulfuric acid or the phosphate buffer solutions. Generally, the maxima could be suppressed readily by addition of a few drops of 0.1% gelatin solution or 0.4% methyl red t o the peroxide solution. Addition of these had little effect on the diffusion current for the peroxides. Effect of Concentration on Polarograms. Several of the hydroperoxides 1% ere investigated extensively for the effect of concentration on the polarographic waves. I n every case, the half-wave potential was independent of concentration. Moreover, the diffusion currents were found to increase linearly with concentration. This is illustrated in Table I11 for three of the hydroperoxides. The markedly smaller diffusion current constant, zd/C, shown for the 1-octyl hydroperoxide, results in part from the high alcohol concentration used with this compound. This effect was mentioned earlier.
6C
50 m
a a
5 40 0
+ 30 a V
20
10
00 ‘0I
-01
-03
-05
-07
APPLIED POTENTIAL, VOLTS vs S
Figure 3.
-09
-11
-1.3
CE
Polarograms for normal hydroperoxides
Solutions approximately 1 X l O - 3 F in hydroperoxides, 0.11‘ in sulfuric acid, and 20% ethyl alcohol
828
ANALYTICAL CHEMISTRY
Electrode Reaction. It seems probable that the electrode reaction involves reduction of the peroxides to the corresponding alcohols as follows: ROOH
+ 2 H T + 2e
+
ROH
+ H20
A calculation of n, the number of electrons involved in the reduction of 1-butyl hydroperoxide from aqueous solutions was made by means of the IlkoviE equation. This yielded a value of approximately 2. I n this calculation a value of 0.77 X cm.* see.-' for the diffusion coefficient of the peroxide was used. This is the value found in the literature for n-butyl alcohol ( 3 ) and i t seems probable t h a t the diffusion coefficient for the corresponding peroxide should be rather close to this. With this approximation n Tvas found to be equal to 1.9. ACKNOWLEDGM EYT
The authors wish to thank Harry S. Mosher and Homer Williams who supplied some of the compounds used in this work. They also wish to thank California Research Corp., whose financial support made part of this work possible.
LITERATURE CITED
(1) Bruschweiler, H., Minkoff, G. J.. -4nal. Chim. Acta 12, 186 (1955). (2) Dobrimskaya, .I.A., Xieman, AI. B., Bcta Physicochim. r.R.S.8. 10, 297 (1939). (3) "International Critical Tables," vol. 5, p. 70, ~\lcGraw-Hill, Xew York, 1929. (4) Lingane, J. J., Laitinen, H. d.,IXD. ESG. CHEM.,ASAL. ED. 11, 504 (1939). ( 5 ) AIacNevin, W.lI.,Urone, P. F.. .ISAL. CHEM.25, 1760 (1953). (6) Xieman, 11. B., Gerber, AI. I., Zhur. Anal. Khim. 1, 211 (1946). (7) Roberts, E. R., Meek, J. S., Analyst 77,43 (1952). (8) Shreve, 0. D., Rlarkham, E. C., J . Am. Chem. SOC.71, 2993 (1949). (9) Shtern, V., Pollyak. S., Acta Physicochim. l'.R.S.S. 11, 797 (1939); J . Gen. Chem., U.S.S.R. 10,21 (1940). (10) Williams, H. R., Masher. H. S., J . Am. Chem. SOC.75, 2984, 2987, 3495 (1954). (11) Willits, C., Ricciuti. C., Knight, H. B., Swern, D., AXAL.C H E Y . 24, 785 (1952). RECEIVED for review October 12, 1955. Accepted January 20, 1956. Division of rinalytical Chemistry, 128th Meeting, ACS, Minneapolis, Rlinn., September 1955.
Voltammetry at Solid Electrodes Anodic Polarography of the Phenylenediamines RONALD E. PARKER' and RALPH N. ADAMS2 Department o f Chemistry, Princeton University, Princeton,
The anodic polarography of the phenylenediamines at a rotating platinum electrode has been examined using a current-scanning technique. 3Ieasurements were made of the variation of Ei/%w-ith pH, and from these data the pK, values of reduced and oxidized species were determined. In all cases examined, linear limiting current 2's. concentration hehaFior was obser, ed. The quantitative determination of 0- and p-phenj lenediamine can be accomplished with a reproducibility of 2 to 394 in the concentration range from to 10-5M, but analysis of mixtures of the two cannot be done except in favorable cases. This study was undertaken to illustrate a potentially valuable adjunct of solid electrode polarography-the measurement of formal potentials of labile organic redox sjstems which cannot be determined by classical potentiometric techniques.
T
H E solid electrode polarography of a variety of organic compounds has been investigated to date. The majority of these studies have concentrated on quantitative analysis by means of limiting current measurements. The important variation of Ell2with p H t h a t accompanies organic polarographic reactions has received lees emphasis. Hedenburg and Freiser examined the mechanism and the dependence of Eli2 us. p H for the oxidation of phenol at a stationary platinum electrode (9). A comprehensive study of the oxidation of phenolic compounds a t graphite electrodes by Gaylor, Elving, and Conrad includes some data on the behavior of E I 2us. pH for phenol, hydroquinone, and 2,4-dirnethyl-G-tert-butylphenol( 8 ) . The \vork reported herein is mainly concerned with the evaluation of dissociation constants of the electroactive species from E , 2's. p H curves taken a t the rotating platinum electrode Although the dropping mercury electrode has been widely used Present address, Johns Hopkins Medical School, Baltimore. J I d . Present address, Department of Chemistry, University of Kansas, Lawrence, Kan. 2
N. J. for this purpose, similar studies with solid electrodes have apparently not, been made. 911 of the present work was done with the rotating platinum electrode using the current-scanning technique previously described ( I , 6 ) . Similar results can be secured using conventional voltage-scan polarography, but it is believed that more reproducible results were obtained with less effort using t,he new technique. 0- and p-phenylenediamine were chosen for study because the diamine-diimine couple represents a highly unstable organic redox system. Comparatively little work has been reported on the anodic oxidation of phenylenediamines a t solid electrodes. I n a comparison of gold, graphik, and platinum electrodes, Lord and Rogers briefly investigated the oxidation naves of 0-, m-) and p-phenylenediamine (11). More attention has recently been given the substituted p-phenylenediamines which are useful as antioxidants and photographic developers. A summary of half-wave potentials of some 50 p-amino-S-dialkylanilines obtained by automatic recording a t a stationary platinum electrode has been given ( 2 , IO). Gaylor, Conrad, and Lander1 recently described a wax-impregnated graphite electrode for the determination of antioxidants of the p-phenylenediamine class (7). EXPERIMENTAL
Reagents. Samples of 0- and p-phenylenediamine were Eastman White Label, used without further purification because it was desired to work with practical samples. Stock solutions (0.01M) for polarographic work Lwre prepared by dissolving the required weight of compound in the minimal quantity of either 1.11 hydrochloric or 1,M acetic acid (for buffer studies) and diluting to volume with air-free distilled water. These solutions were prepared fresh each day to minimize the effect of air oxidation. Britton and Robinson buffer solutions were used for the p H studies and were prepared according to the directions given by hliiller ( 1 5 ) . The stock buffer solution prepared was 0.0411.1 in phosphoric, boric, and acetic acids, respectively. Varying volumes of this solution were mixed with 0.2M sodium hydroxide to prepare the desired buffers. The exact, p H was measured at the end of each polarographic run using a Leeds I%Xorthrup p H meter, which was standardized prior to each of these p H deter-
