1078
ANALYTICAL CHEMISTRY
the pattern of the rising boundaries is slightly greater than that from the other side of the channel. This doubtless results from incomplete “drainage” during shifting of the initial boundaries into view. Thus contamination with protein on the side in which the boundary descends raises the index of the buffer solution more than contamination by buffer lowers the index of the protein solution on the other side. The 2-ml. cell has also been used with the autocollimating camera for diffusion measurements. In the case of 0.621y0 dlalanine diffusing into water a t 1 O , a sharpening flow of 1 mi. per minute for 30 minutes was used, the outside diameter of the capil1:iry siphon being 0.65 mm. The pattern of Figure 5, a, was recorded during this process while those at b and c were obtained after diffusion for 500 and 1000 seconds, respectively. Here the total number of fringes in the pattern is 40.85 and the average deviation, from the mean value of 4.320 X of the coefficient from each of the ten exposures that were made during the diffusion period was 0.05%. The same solution in the 11-ml. cell gnve 4.317 X (6). ACKNOWLEDGMENT
The developments reported here would not have been possible
without the care and ingenuity exercised by Emil Maier of the Pyrocell Manufacturing Co. in the construction of the new cells. The author is also indebted to R. J. Slater of the Hospital of the Rockefeller Institute for much interesting material, both normal and pathological, that has been used in testing the apparatus and to D. A. JIacInnes of these laboratories for his criticism of this manuscript. LITERATURE CITED
(1) Antweiler, H. J., Kolloid-Z., 115, 130 (1949). (2) Chambers, L. A , private communication; U. S. Patent 2,412,602 (1946). (3) Labhart, H., and Staub, H., Helv.Chiin. Acta, 30, 1954 (1947). , 346 (1951). (4) Longsworth, L. G., -4s.4~.C H E M .23, (5) Longsworth, L. G., Chem. Revs., 30, 323 (1942). (6) Longsworth, L. G., IND. ENG. CHEW, d s a ~ ED., . 18, 219 (1946). -, (7) Longsworth, L. G., J . Am. Chem. SOC.,6 5 , 1755 (1943). (8) Ibzd., 74, 4155 (1952). (9) Longsworth, L. G., ”IIethods in Medical Research,” Vol. 5. p. 63, Chicago, Ill., Year Book Publishers, Inc., 1952. (10) Moore D. H., and Khite, J. H., Rev. Sci. Instr., 19, 700 (1948). (11) Svensson, Harry, and Odengrim, Karl, Acta Chem. Scand., 6,720 (1952). ~
RECEIVED for review January 26, 1953. Accepted M a r c h 23, 1953.
Polarographic Oxidation of Phenolic Compounds V. FRANCES GAYLOR, Chemical and Physical Research Division, The Standard Oil Co. (Ohio), Cleveland, Ohio, PHILIP J. ELVING, University of Michigan, Ann Arbor, Mich., AND ANNE L. CONRAD, Chemical and Physical Research Division, The Standard Oil Co. (Ohio), Cleveland, Ohio Because of the limitation of the dropping mercury electrode in reference to the polarographic investigation o€oxidation reactions, the use of solid electrodes, especially with organic compounds, seems of great potential value. A polarographic method, using an indicator graphite electrode, has been developed for studying the oxidation of organic compounds and has been applied to the study of phenolic compounds. Reproducible results can be easily obtained for a single compound; results are comparable for a series of compounds if predetermined conditions are carefully controlled. Variables affecting the reproducibility of the method have been explored, and the basic information and experimental conditions necessary for the quantitative determination of compounds by the method have been established. The relative ease of oxidation of compounds can be rapidly measured. The behavior of three typical phenolic compounds (phenol, hydroquinone, and 2,4-dimethyl-6-tert-butylphenol) has been investigated in detail.
A
METHOD for measuring the relative stability or oxidiaability of organic compounds would be advantageous for many purposes, I n most cases, these compounds oxidize irreversibly and their oxidation potentials cannot be determined by standard methods. Fieser (2) has measured oxidizability of organic compounds by determining the critical oxidation potential: this necessitated a long and tedious procedure. It was desired to develop a more rapid method, for which purpose the polarographic technique seemed especially well suited. Conventional dropping mercury electrodes could not be used at potentials positive enough t o oxidize most organic compounds, owing to prior oxidation of the mercury a t potentials less positive than 0.4 volt. This necessitated the use of solid electrodes. Rogers, Miller, Goodrich, and Stehney ( I O ) reported the use of solid electrodes in inorganic work and proved that automatic
recording would give satisfactory results. Others ( 1 , 5, 4, 15’) have reported studies of the oxidation of organic compounds using platinum electrodes. In using any solid electrode with a nonrenewable surface, previous history of the electrode may affect, results. Therefore, the necessity of maintaining constant surface conditions became an important factor in securing reliable data. Rogers and Lord (9) reported that a graphite electrode could be used for obtaining oxidation waves of organic compounds; a new and reproducible surface was secuied for each run by breaking off the tip of the electrode. This offered a definite advantage over the platinum electrode. The work described in this paper was based upon their contribution. The method discussed employed wax-coated graphite electrodes, to which a positive potential was applied us. the standard saturated calomel reference half-cell. Buffered alcohol-water solutions of the compounds were used. Polarograms were automatically recorded after degassing of the cell test solution and preconditioning of the electrode. Phenolic-type compounds were selected for study because of their wide utility as antioxidants and because they had been previously investigated potentiometrically (i?) and, in the case of hydroquinone, polarographically ( 7 , 8). The specific compounds investigated were phenol, the first member of the series, hydroquinone, which exhibits reversible oxidation-reduction, and 2,4-dimethyl-6-tert-butylphenol, which has been highly recommended as an oxidation inhibitor (11 1.
APPARATUS
Current-potential curves were recorded by a Sargent Model XXI polarograph; potential measurements were checked with a potentiometer. A saturated calomel electrode was used cathodically as an external reference electrode and was connected to the polarographic cell through an agar bridge. The cell (25-ml. capacity) was jacketed with provision for circulation of water, controlled a t 20” 5 0.1 O C. Cell resistances were measured with a conductivity bridge (Industrial Instruments Co., Model RC-
BC).
The electrode was a 0.25-inch graphite rod of high purity. The first electrode conslsted of spectrographic grade A graphlte elec-
V O L U M E 25, NO. 7, J U L Y 1 9 5 3
1079
trodes manufactured by United Carbon Products Co. ; the second electrode consisted of special graphite spectroscopic electrodes manufactured by National Carbon Co. Cold electrodes were brushed with melted ceresin wax, so that the surface was uniformly coated. Care was taken to prevent the impregnation of the rod with the wax; otherwise, erratic polarographic curves resulted. The lower tip of the electrode, about 0.25 inch, vlas then broken off for insertion into the cell before each run. A short piece of rubber tubing containing mercury served as a contact between the graphite electrode and the lead to the polarogiaph. REAGENTS
The chemicals used were hydroquinone (Eastmari Kodak white label grade), phenol (Baker c.P., purified by distillation), 2,4-dimethyl-6-tert-butylphenol (complimentary sample from the Shell Development Co.), and isopropyl alcohol (Fisher c.P.). The dimethylbutylphenol (DMBP) was purified by fractional crystallization in the Firestone Research Laboratory through the courtesy of G. E. P. Smith, Jr., and L. J. Kitchen. The purified compound had a purity, based on melting point behavior, of 97 mole yo; its refractive index checked the literature value.
1.0
0 33 rn&
mM
/
PHENOL
DISCUSSION OF METHOD
Typical polarograms obtained from the three compounds studied are shown in Figure 1. In the case of hydroquinone id was measured from point A to the residual current. The decrease of current a t B was similar to that obtained with the dropping mercury electrode. This regressive behavior has been ascribed to a coating of merdury phosphate on the dropping mercury electrode ( 6 ) ; its presence in the absence of mercury and phosphate ion indicates that the cause is a more basic phenomenon. Small maxima sometimes appeared a t higher conrentrations of phenol, but these were usually not large enough to mask the limiting current. The phenol wave illustrates the magnitude of the residual current sometimes obtained and the need for blank correction in such case?. Dimethylbutylphenol gave three waves a t pH 5.2. The use of a 50% alcohol solution as solvent may have given pH values which were only approximate and potential values which contained small errors due to liquid junction potential effects. This would in no way affect the general conclusions drawn but only calculations based on the test solution pH. Anodic preconditioning of the electrode before each run greatly improved the reproducibility of the results. Other methods of preconditioning may prove to be similarly satisfactory. For example, preconditioning the electrode in the absence of sample gave results similar to those described. The variables which were explored are discussed below.
HYDROQUINONE
A
0
to 2
toe
104
+o e
VOLTS
Figure 1. Polarograms of Compounds Studied pH 5.2.
Graphite us. S.C.E. 1.24 m v . per second. corrections
No i R
Aqueous buffer stock solutions aere prepared by adding acid or base to a solution of 1 mole of the major buffer component until the pH was such that, on subsequent dilution with an equal volume of alcohol, the pH was that desired; the final volume of the aqueous stock solution was 1 liter. Buffer components were for p H 1.2 potassium chloride-hydrochloric acid; for pH 5 . 2 sodium acetate-acetic acid; and for pH 8.2 ammonium chloride-animonium hydroxide. Other buffers \+-ere the usual 3\IarIlvaine, and Clark and Lubs types. Commercially available nitrogen was purified and conditioned by passage through alkaline pyrognllate, water, and a portion of of the test qolution. PROCEDURE
The cell test solution was prepared by mixing an alcoholic solution of the sample with the desired buffer. In a 50-ml. volumetric flask, 25 ml. of buffer solution plus the desired aliquot of the sample solution 1%as placed; the final volume was adjusted with isopropyl alcohol. Twenty-five milliliters of this solution was placed in the prebubbler in the nitrogen dream and 25 ml. in the cell. Kitrogen was passed through the cell solution for 15 minutes, during which time anodic pretreatment of the electrode was conducted. This consisted first of a 10-minute application of a constant potential a t some point on the diffusion plateau of the sample; for comparative purposes, a potential was taken equal t o 90% of the total potential to be applied. Following this, the electrode was allowed to equilibrate a t the initial potential for 5 minutes. After the nitrogen stream was transferred to the cell atmosphere, the anodic wave was recorded in a positive direction. At the completion of the run, the polarographic leads were disconnected and the resistance of the cell was measured. Polarograms were corrected for iR drop across the cell. The cell R did not exceed 700 ohms and the zR correction was less than 0.05 volt in every case. In the majority of cases a blank correction for residual current was also necessary. The half-wave potentials, El/z,and diffusion cwrents, id, were measured by geometrical methods.
2
1
0
20
40
062
Figure 2.
rn:%pE~
60
80
sEC
Effect of Polarization Rate on HalfWave Potential
Effect of Polarization Rate on Half-Wave Potentials. A change in the rate of applying the potential may shift the position of the wave on the potential scale. More positive values for E112 were obtained with faster polarization rates, as illustrated in Figure 2, which compares El/2 a t three different polarization rates under otherwise comparable conditions, The data in this graph are comprised of values for all three compounds measured a t various pH levels. Below 0.3 volt, changes in polarization rate caused no significant difference in Ellz. Above 0.3 volt, however, increasing the rate of polarization caused large positive shifts in E112. For example, a wave with an apparent E112 of 0.60 volt a t a polarization rate of 0.62 mv. per second would appear to have E1/2 a t 0.68 and 0.79 volt for rates of 0.93 and 1.24 mv. per second, respectively. Thus in order to obtain comparable data on El, t measurements, identical polarization rates should be used. However, electrodes from different sources used a t identical polarization rates may not give the same values for El/n. This has been found to be due to a difference in magnitude of the shift of E l / ?with polarization rate for different electrodes.
ANALYTICAL CHEMISTRY
1080
It is possible that the shift of with polarization rate could be explained by the failure of the electrode to assume the actual value of the limiting current corresponding to the potential of the electrode as rapidly as the potential is applied. Increasing the rate of polarization would increase the current lag, with a concomitant shift in E l l 2 to more positive values. This current behavior would be in conformity with the known behavior of stationary solid electrodes and would be due to failure of the concentration gradient to be set up instantaneously (6). The slopes of the waves are in line with this statement. At slower rates of applied potential, the waves were steeper, giving less positive values for El,2 . ,d AT 0 6 2 m v / S E C 'd AT I 2 4 m v / S E C
f'
I N
= L
N
o
+O.SOk.,
N
I,
'I, N
PH
Figure 4. HYDROQUINONE
PHENOL
1.24 my. per second. 0.33 m M
OVBP
Figure 3. Effect of Polarization Rate on Diffusion Currents
Effect of Polarization Rate on Diffusion Currents.
Effect of pH on Half-Wave Potential
Values of
id were also found to be dependent upon the rate a t which the potential was applied. Faster polarization rates resulted in larger values for id, Figure 3. The current, reported in microamperes per millimole, was an average value calculated from the various concentrations run a t three pH values. The average effect of increasing the polarization rate of 0.62 to 1.24 mv. per second was to increase id by 40 to 50%. Within the limits of experimental error, this figure remained constant through a series of 65 runs in which d8erent concentrations, pH values, and electrodes were used. Thus it can be seen that polarization rate must be controlled if comparable current values are to be obtained. The enhanced eurrent a t higher polarization rates may be due to nonequilibrium conditions or to stirring effects, resulting from the density gradient caused by depletion a t the electrode surface. Effect of Polarization Rate on Wave Slopes. The n values as calculated from the wave slopes have been interpreted in the case of reversible reactions as being equivalent to the number of electrons involved in the potential-controlling step of the electrode reaction, Approximate n values were calculated from the relation : E3/4 - Ella = 0.055/n (at 20' C.) which asmmes a straight line between the potential limits equivalent to 0.25 and 0.75 id. The n values varied with polarization rate; the average effect of increasing the polarization rate from 0.62 to 1.24 mv. per second was to decrease the n value by 1.0 unit. This is in accordance with the concept of current lag on the electrode discussed previously. Comparison of Electrodes. Oxidation waves obtained with electrodes from two different sources were compared a t identical polarization rates. At a rate of 1.24 mv. per second the 21 waves measured were found to differ only within limits of experimental error. For either type of electrode the standard error in the determination of E1/2 was 0.02 volt and the error in i d was 20% of the current level. At lower polarization rates, poorer agreement between the two electrodes was obtained. The fact that both E112and wave slopes varied between electrodes was probably due to the differences in magnitude of the current lag. Effect of pH. Values of Ell2 decreased aTith increasing p H in the range of pH 1 to 10. The experimental values of Ell2for
hydroquinone and phenol measured with a single electrode type are plotted against the corresponding pH values in Figure 4. These data may be related to the standard oxidation potential, EO,by means of the equation: Ed
= E,/,
+ 0.058 pH (at 20" C.)
Conversion of the experimental E l / *gave average values of Eo of 0.454 and 0.959 volt us. the saturated calomel electrode for hydrcquinone and phenol, respectively. These experimental Eo values were in excellent agreement with the 0.453 and 0.979 volt us. the maturated calomel electrode reported in the literature for the two compounds. [Eo for phenol was calculated from the critical oxidation potential reported by Fieser ( 2 ) . ]
1.24 m v . per second
The id of hydroquinone was independent of p H (Figure 5), while the id of phenol was greatest a t pH 1.2 and decreased with subsequent increase in pH. (The sizes of the points in Figures 4, 5, and 6 are proportional to the standard error, determined from a large number of runs.) In the case of dimethylbutylphenol (Figures 4 and 6 ) the number of oxidation waves obtained was determined by the p H of the cell medium. At pH 5.2 three well-defined waves were obtained. The id of the wave of lowest El/*was approximately equal to the sum of the id of the other two wavm. El/%values
V O L U M E 25, NO. 7, J U L Y 1 9 5 3
1081
were greatest at p H 1.2. However, in this highly acid medium the first two waves were no longer well separated, as Ella of the first wave increased more than that of the second wave. The id of the first wave was greater a t pH 1.2 than a t 5.2, in a manner analogous to that of the phenol wave. The second wave showed no measurable change in id, whereas the highest wave showed a great increase in id a t pH 1.2. At p H 8.2 E1/z values were lower than those obtained a t pH 5.2. The two lower waves were well defined with little or no change in limiting current. The third wave was so small as to be discernible only a t high concentrations of the compound. d t pH 10 only one well-defined wave a t a very low potential was obtained. The waves for hydroquinone and dimethylbutylphenol very much depressed at pH 3'29 gave an apparently normal wave. The decreased current may be the result of the interaction Of the with the buffer components (phosphate and citrate) to produce species which diffuse more slowly. Table I.
Shift of Half-Wave Potential with pH
Compound
Polarization Rate, Mv./Second 0 62 0 93 1 24
p H 1.2 t o 5.2 (Calcd. Shift = 0.232) Hydroquinone 0.17 0.18 0.15 0.18 Phenol 0.14 0.18 D M B P , wave 1 D M B P , wave 3 0.14 0.19 pH 5.2 t o 8.2 (Calcd. Shift = 0.174) Hydroquinone 0.16 0.16 Phenol 0.11 0.14 D M B P , wave 1 0.13 0.13 D M B P , ware 2 0.13 0.15
0.20 0.22
0.21 0.23
0.17 0.16 0.17 0.21
Depression of the id of dimethylbutylphenolofferedan explanation for the absence of the second wave of this compound at pH 3.2. il large depression of its normal wave height resulted in a very small wave not discernible a t the sensitivity level used. k o attempt has been made to assign the wakes observed for the three compounds studied to definite chemical reactions. However, a smooth wave was obtained for quinhydrone a t the graphite electrode, which seemed to indicate that the hydroquinone wave was due to the normal two-electron oxidation to quinone. Based on this assumption, the phenol wave a t pH 1.2 was probably the result of a four-electron process, since the average id of phenol a t pH 1.2 was approximately twice that of hydroquinone. At other pH values the average id of the two waves were approximately equal, indicating a two-electron ouidation of phenol. rlt pH 1.2 the sum of the three waves of dimethylbutylphenol indicated a collective four-electron process, while the largest wave a t pH 5.2 represented a two-electron process; a t other pH values, disappearance of some of the waves prevented more definite statements. I t may be significant that the height of the dimethylbutylphenol waves a t pH 5.2 had the ratio of 2 : l : 1. The enhanced currents a t p H 1.2 for phenol and dimethylbutylphenol may have a common cause. The foregoing data were obtained a t one polarization rate and with a single type of electrode. The effect of p H on Ell? may be influenced by polarization rate, since the magnitude of the effect of polarization rate on potential varied with the potential itself. The magnitude of the latter shift therefore depended upon the position of the wave on the potential scale, being greater a t more positive E l l Bvalues. Consequently, the apparent shift due t o pH was actually the sum of the p H effect and of the change in magnitude of the polarization rate effect in going from one region of potential to another. The total observed effect has been summarized in Table I where the shifts are tabulated for three polarization rates. Comparison of results from the two electrodes gave agreement within the limits of experimental error.
Table 11.
Effect of Temperature on Diffusion Current
(pH 5.2, 1.24 mv. per second, 2 concentration levels) Temp. Coeff., Compound i d a t 20' C. i d a t 30' c. % Phenol 6.9 7.2 0.4 4.8 5.2 0.8 DMBP. wave 1 8.1
8.7
0.7
DRIBP, wave 2
2.4 4.2
3 0 5.4
2.5 2.9
DLIBP, wave 3
2.0
3.0 6.9
5.0 5.3
4.5
Effect of Temperature. Currents were found to increase with increasing temperature. Because adsorption-controlled waves normally show the reverse effect, none of the waves studied could be classified as being due to adsorption. Temperature of id were calculated for each wave from measurements a t 20' and 30' C. (Table 11). Phenol and the first wave of dimethylbutylphenol showed similar temperature coefficients, 0.4 and O.8y0 per O C., respectively. These values were sufficiently low to indicate that the limiting currents were probably largely diffusion-controlled. The second wave of dimethylbutylphenol gave a higher value of 2.7% but might still be diffusion-controlled. The value obtained for the third wave, 5.2%, was sufficiently high to be questionable as to its kinetic, diffusion, or composite origin. X'o appreciable differences in E l / *of the waves were found a t the two temperature levels. Effect of Concentration. Over the concentration range from 0.1 to 1.0 millimolar, a linear relationship was found between concentration and i d for each of the three compounds. The tests were carried out in solution buffered a t pH 5.2 and a t a polarization rate of 1.24 mv. per second. The straight-line relationship between id and concentration is the basis for using the polarographic method for quantitative determinations of the compounds. 30
r I---_
',
--____
I---
WAVE
I
WAVE
3
- - - - - - - ------- - - - _ _ _ _ _ _
I
O
12
3.2
52
82
IO 0
PH
Figure 6.
Effect of pH on Diffusion Current
Dimethylbutylphenol.
1.24 m v . per second
In most reversible reactions a t a dropping mercury electrode, Ells has been found to be independent of concentration of the active component. Reversible organic redox systems as well a8 many irreversible organic reactions may also show a similar behavior. A4smeasured with the graphite electrode, Ella of hydroquinone did not shift with concentration variation, which, in conjunction with the other evidence, would tend to indicate that the reaction was reversible a t the solid graphite electrode. The smooth wave obtained for quinhydrone a t the graphite electrode substantiates the reversibility of the reaction. El/Z values of the waves of phenol and dimethylbutylphenol were dependent upon the concentration. With tenfold dilution of these compounds, Ellz shifted in a positive direction. The average differences of E l / % obtained on three or more sets of runs
1082
ANALYTICAL CHEMISTRY
are tabulated in Table 111. These shifts indicated irreversible electrode reactions or deposition on the electrode if these results could be considered analogous to those of the dropping mercury electrode system. Interpretation of these shifts in terms of number of electrons involved in the reactions may not be justified a t the present time.
Table 111. Effect of Tenfold Dilution of Oxidizable Component on Half-Wave Potential Compound Hydroquinone Phenol DMBP, wave 1 D M B P , wave 2 DMBP, weve 3
(1.0 and 0.1 m M ) PH 1.2 5 2 8 2 5 2 8 2 1.2 5.2 8.2 5.2 8.2 5.2
Average Shift in E l / l KOchange KOchange No change 0 022 0.038 0.016 0.026 0.035 0.022 0.034 0,026
In addition to the change in El,2 with dilution, wave slopes, expressed as n values, also tended to become larger a t higher dilution. CONCLUSIONS
-4 polarographic method has been developed for studying the oxidation of phenolic compounds, which should be equally applicable to the investigation of other types of compounds, Comparable data on the ease of oxidation of a large number of compounds can be rapidly obtained. Quantitative measurement of their concentration can also be made. Standardization of experimental conditions is a primary requirement for all comparative and quantitative work. Variation in polarization rate, pH, temperature, and concentration may
affect either the half-wave potential, the diffusion current, or both. However, these variables can be easily controlled. The reIative ease of oxidation of a large group of phenolic-type compounds has been measured by the technique described. ACKNOWLEDGMENT
The work presented in this paper was performed in the laboratories of the Standard Oil Co. of Ohio, and the authors wish to express appreciation for permission to publish the material. LITERATURE CITED
Bent, R. L., Dessloch, J. C., Duennebier, F. C., Fassett, D. W., Glass, D. B., James, T. H., Julian, D. B., Ruby, W. R., Snell, J. M.,Sterner, J. H., Thirtle, J. R., Vittum, P. W., and Weissberger, A , , J . Am. Chem. Soc., 73, 3100 (1951). Fieser, L. F., Ibid.,52, 5204 (1930). Hedenburg, J . F., and Freiser, H., Pittsburgh Conference on Analytical Chemistry, March 1952. Julian, D. B , and Ruby, W. R., J . Am. Chem. Soc., 72, 4719 (1950).
Kolthoff, I. M., and Furman, K. H., “Potentiometric Titrations,” 2nd ed., New York, John Wiley & Sons. 1931. Kolthoff, I. M., and Lingane, J. J., “Polarography,” 1st ed., New York, Interscience Publishers, 1941. Muller, 0. H., J . Am. Chem. Soc., 72, 4719 (1950). Muller, 0. H., and Baumberger, J. P., Trans.Am. Electrochem. Soc.. 71. 169. 181 (1937).
Rogers, L:B., Lnd Lord, S. S..Pittsburgh Conference on -4nalytical Chemistry, March 1952. Rogers, L. B., MiIIer, H. H., Goodrich, R. B., and Stehney, -4.F., ANAL.CHEM.,21, 777 (1949). Rosenwald, R. H., Hoatson, J. R., and Chenicek, J. A., Ind. Ena. Chem.. 42. 162 11950).
G., and Atamanenko, N. N., Zawdskaya Lab., 15,
(12) Skobets, E. 1291 (1949).
RECElrED for review October 20. 1952. Accepted April 3, 1953. Presented before the Division of Analytical Chemistry a t the 122nd Meeting of the A h r E R r C A N CHBMICAL SOCIETY, Atlantic City, N. J. 9 portion of t h e reported work has been submitted by V. F. Gaylor in partial fulfillment of the requirements for the M.S. degree a t Western Reserve University.
Polarographic Analysis of Mixtures of Fumaric and Maleic Acids, and Their Diethyl Esters PHILIP J. ELVING’, AARON J. MARTIN, AND ISADORE ROSENTHAL* The Pennsylvania State College, State College, Pa.
T
H E polarographic method for the simultaneous determination of maleic and fumaric acids is especially useful since it is difficult to distinguish these stereoisomers by the usual methods for determining acids ( 5 ) . Apparently no method has been described for the analysis of mixtures of these acids with their diethyl esters. Such an analytical method would be of advantage for following saponification and esterification reactions, or for the detection and determination of the compounds as impurities or as products of side reactions. The polarographic waves most suitable for the simultaneous determination of maleic and fumaric acids for analysis are obtained a t p H 8.2 in ammonia buffer ( 5 ) . The method is applicable for concentrations up to 2.6 m M with an accuracy of 2 to 3 relative yo. The potential-spot current-reading technique can be used successfully. Polarographic data for the diethyl esters ( 2 , s )show that, below p H 5, the reduction wave of diethyl maleate is sufficiently separated from the wave for diethyl fumarate t o permit simultaneous analysis of mixtures of these compounds. 1
3
Present address, University of Michigan, Ann .4rbor, hlicb. Present address. Rohm & Haas Co., Philadelphia, Pa.
The mechanism of the polarographic reduction of unsaturated acids such as maleic and fumaric acids has been discussed by many workers (cf. reference 5 ) . Most of this work was done in acid solution; some of the work in neutra1 and alkaline solution is of doubtful validity because of the use of unbuffered solutions. More recent work (3) includes a study of the acids in adequately buffered basic solutions. A Leeds and Northrup Model E Electro-Chemograph was used. The dropping mercury electrode (Corning marine barometer tubing) had a capillary constant of 1.65 mg.*/Ssec.-*/2a t 75 cm. into distilled water, open circuit, a t 25.0” C. This moderately large head of mercury was used to increase the diffusion current so that lower concentrations could be measured with greater accuracy. The polarographic cell ( 4 ) was maintained at a temperature of 25.0” += 0.1” C. The nitrogen used for deoxygenation of all test solutions was purified and conditioned ( 3 ) . The mercury was purified according to Wichers (6). All pH measurements were made on a Beckman Model G pH meter. Resistances as measured on a General Radio Model 650-A impedance bridge were found to be negligible. Maleic and fumaric acids were prepared and purified as described (1). Diethyl fumarate and diethyl maleate (Carbide and Carbon Chemicals Co.) were used as received; boiling points and polarographic analysis indicated their purity to be adequate. The purity of the C.S.P. grade 95%
