Anodic Voltammetry of Phenols JOHN F. HEDENBURG AND HENRY FREISER Department of Chemistry, University of Pittsburgh, Pittsburgh, Pa. This study was undertaken to investigate the applicability of anodic voltammetry to phenols, in the hope that a rapid quantitative method of phenol analysis could be developed. This work was also motivated by the desire to learn more about the oxidation of phenols, with a view to developing a method for evaluation of phenols as antioxidants. It has been found that phenol can be oxidized at a micro platinum anode, yielding a current-voltage curve similar to the conventional polarographic wave. A reproducible half-wave potential, shown to be a function of pH, is obtained. At pH 9, the halG
T
HE object of this study was to investigate the voltammetric behavior of phenols, in order to develop a quantitative method of phenol analysis, to obtain a measure of the ease of oxidation of phenols, and to develop a method for evaluating phenols as antioxidants. As phenols cannot be oxidized reversibly, they exhibit no thermodynamic oxidation potentials. Two potentiometric methods for studgng phenol oxidation involve the “apparent oxidation potential” ( 5 ) and the “critical oxidation potential” (6). The procedures for obtaining these values are lengthy and the results can be considered only approximate, since theoretical objections to both can be raised. Muller ( l a ) has shown that the apparent reduction potential of a compound corresponds to the decomposition potential on the polarographic wave of the compound. EXPERIT1E 3 T A L
Reagents. Reagent grade phenol was used. The other phenols used in this work were furnished by R. S. Bowmaii ( I ) , and were used without further purification. Physical Constants of Phenols Phenol o-tert-But,ylphenol m-lert-Butylphenol p-tert-Butylphenol
R P., 20 hfm.,
O
C
113.0 129.5 130.0
SI.P., O
c.
...
43 100
A 0.1 ill stock solution of phenol in water was prepared and used to prepare the more dilute solutions. Solutions of the tertbutylphenols were prepared in the same manner, using (1 to 11 dioxane-water as the solvent. Buffer solutions were prepared by dissolving the buffering agents, reagent grade (disodium phosphate for pH 9.0 and sodium hydroxide-boric acid for p H 10.4), in 0.1 A- potassium nitrate solution and adjusting to the desired pH value with the aid of n pH meter. Total conrentration of buffering agent was approximately 0.1 X . In certain experiments Coleman buffer tabletb were dissolved in 0.1 N potassium nitrate solution to prepare the buffer solution. Sample< vere prepared by pipetting from 1 to 5 nil. of phenol solution to give 100 ml. of solution. Apparatus. .I Fisher Elecdropode was used in the early part of this study, and a Leeds and Korthrup Electro-Cheinograph Model E was used in the latter part of the work. All data obtained with this instrument were taken a t 25’ j=0.5’ C. A very simple cell, either a 100-ml. beaker or a Petri dish, was found suitable for the polarographic experiments. A Beckmnn pH meter, Model €12, was used in all p€I measurements. RESULTS OF EXPERIMENTS
Use of Dropping Mercury Electrode. An attempt to study the oxidation of phenol was first made with a n anodically polarized dropping mercury electrode using the Elecdropode. While results obtained with phenol were different from those obtained with blanks, it was impossible t o make analytical use of the
wave potential has a value of 0.52 v o l t . i n equation has been derived which closely describes the pH dependence of the half-wave potential. A plot of log I / ( h - I ) us. E gives a straight line, indicating a polarographically reversible one-electron oxidation to the phenoxide free radical. The diffiision current is found to be closely proportional to the concentration and essentially independent of the pH. This new method of analysis for phenols will also permit an electromotive series to he constructed for phenols and enols whose over-all oxidation is irreversible,
data. This &-asprobably due to the fact that the range of potential in which mercury oxidizes overlaps that in which the phenol wave occurs. When it was learned that Julian and Ruby (9) had successfully determined half-wave potentials of irreversibly oxidized systems using a platinum microelectrode, it was felt that this technique might be applicable to the study of phenol oxidation. 9 n electrode similar to that ujed by those investigators was prepared from an 18-mil platinum wire 7.5 mm. long, arranged so that it v a s parallel to the bottom of the cell. A saturated calomel electrode was used as the cathode. The Elecdropode was used t o obtain the current voltage curves. At a pH of 9.0 in a phosphate buffer the curve for 0.01 11‘ phenol departed from that of the blank a t 0.25 volt and rose sharply. Since phenol is only 10% ionized a t 3, p H of 9.0, a sodium hydroxide-boric acid buffer a t B pH of 10.4 was next prepared. -4gain the departure of the phenol curve from the blank was observed. It was found that if a second run was made on the eame sample, a curve identical with that of the blank was obtained. When the potential was raised over 0.5 volt, the current rose and then gradually dropped back. If left a t this potential long enough, the current dropped back to that of the blank curve. A black deposit was found on the electrode after such treatment. This deposit was insoluble in acetone, dioxane, chromic acid, nitric acid, or strong alkali. The formation of the deposit gradually insulated the electrode from the cell, causing the observed current drop. Attempts t o eliminate the formation of the deposit by adding dioxane to the cell and purging with nitrogen failed. The onlv way found to clean the electrode was to burn off the deposit i n the flame of a Bunsen burner for 30 srronds prior to csach run.
M , reducing By lowering the phenol concentration to 1 X the quantity of electricity passing through the cell during a run, and leaving the circuit closed only long enough to obtain a steady reading, curves of the type shown in Figure 1 were produced. -41t,hough the results were improved, the limiting current still could not be evaluated. Further reduction in concentration to 1 X 10-6 ill produced a wave with a definite limiting current, as shown in Figure 1. The voltage a t the beginning of the wave (decomposition potential) increased as concentration decreased. 8 s electrode coating occurs throughout the run, the extent of the coating is a function of the run time. As it was not possible to control the run time using a manual instrument, it was felt that the use of an automatic recording instrument would improve the character of the curves. This would permit minimization of the effect of coating by reducing the run time and by making the run time reproducible. Use of the Electro-Chemograph. The same platinum microanode and calomel cathode were used with the Elecdropode and with the Electro-Chemograph. Excellent waves were obtained, as shown in Figure 2. The method for determining diffusion current and half-wave potential is shown in Figure 2 1355
ANALYTICAL CHEMISTRY
1356 Effect of Concentration. The data showed that diffusion current is dependent on concentration. All values of K I , the proportionality constant between diffusion current and concentraThe average value is 1.4 tion, fall between 1.0 and 2.0 X X 10-2 i 14% (standard deviation = 0.4). Those results are more variable than those usually obtained in conventional polarography.
0
Figure 1. Effect of Phenol Concentration pH 10. Fisher Elecdropode. Full sensitivity. 1
=
0.25 mm.
Reproducibility. HALF-WAVE POTENTIAL. The most reproducible feature of the phenol wave is half-wave potential, which has been found to be 0.51 i0.02 volt a t p H values of 9.0 and 10.4. This is an average of 50 determinations. DIFFUSION CURRENT.The average Kl is 1.4 X lo-* i 14%; hence diffusion current may be expected to be reproducible to within 14% of the mean value expected a t any concentration. At concentrations from 1 X 1O-j to 3 X 10-5 ilf the wave height is small and errors in measurement are magnified. At concentrato 10 X 10-j JI,resin formation on theelections from 8 X trode becomes significant. Best results should be expected beM . This is shown by the two series tween 3 X 10-5 and 8 X of curves, both of which were run Tyithout changing the solution or burning the electrode between runs. There r a s very little change in K1 in the first series at 4 X 10-6M, vhile in the second M , K: dropped considerably, indicating the exseries a t 8 X tensive formation of the electrode deposit. This shows that the condition of the electrode markedly affects the relation of diffusion current to concentration. For this reason, preparation of the electrode prior to a run was kept under close control. Storage of the sample before running does not affect the reproducibility. A 100-ml. sample of 4 x 10-5 icf phenol a t a p H of 9.0 was prepared and divided into two equal portions. One portion was run immediately, while the other portion stood for 0.5 hour and was run with a burned electrode. After standing another 0.5 hour, the same sample was run with a burned electrode. The identical curves show that the sample ~villnot decompose in storage for an hour and that duplicate runs can be made on the same sample. Changing the distance between the microanode and the calomel cathode caused variations in diffusion current and in the general
Table I. Voltammetric Experiments with tert-Butylphenols p H = 9.0. T = 25O C. Concn. = 10 X 10-6 Jf. Scale = 3 pa. El/za0 Id, ED, EM, Compound Volt pa. Volt Volt 0.40 0.58 60 n-tert-Butylphenol 0.47 1.02 61 o-tert-Butylphenol 0.38 0.72 0.31 0.48 0.33 0.47 62 tert-Butylphenol 0.41 0.75 63 .&enol 0.515 1.14 0.42 0 60 0 Et,n is the half-wave potential, I d , the diffusion current, E p , the decomposition potential, and E M , the value of the potential at the hmiting current. KO.
shape of the curve. I t is desirable that the locations of the electrodes be fixed relative to one another. Effect of pH upon D i h s i o n Current and Half-Wave Potential. Curves were run on samples prepared using Coleman buffer tablets. The diffusion current is pH-independent in the range of pH values from 5.0 to 11.0. The half-wave potential decreases rapidly as pH increases up to a p H of 9.0. At higher pH values it is essentially independent of pH. Number of Electrons Transferred. From the slope of the line obtained from the plot of log Z / ( I d - Z) versus E, n, the number of electrons involved in the oxidation, can be calculated in the usual way. The average (over six determinations) number of electrons transferred is 0.996, showing that the oxidation step responsible for the voltammetric wave involves the transfer of one electron. Voltammetric Behavior of Other Phenols. 0 - , m-, and p-tert-butylphenols mere run in the same manner as phenol (Table I). The alkyl substitutent in the meta position has little effect on the ease of oxidation, but the inductive effects of the ortho and para substitution of the alkyl group render these compounds more easily oxidized than phenol. The results indicate that differentiation among phenols is possible by voltammetric studies. A4seach wave covers a potential range of approximately 0.2 volt, simultaneous determination can be carried out only in cases where the half-wave potentials differ by 0.2 volt. Where the differences are less than this amount, the waves should overlap, resulting in one wave, the diffusion current of which would be proportional to total phenols present.
I 14-
12-
lo-
A0
8-
6.
4.
2.
0.
,
2
.3
4
.S VOLTS ys SCE
.6
I
7
.a
Figure 2. Waves Obtained with Electro-Chemograph
V O L U M E 25, NO. 9, S E P T E M B E R 1 9 5 3
1357
Practical Application. -4buffered (Coleman tablet) sample of waste liquor from a coke plant ammonia still was run in the usual manner. The data, recorded in Table 11, show a phenol concentration in the liquor of 666 & 46 p.p.m. The phenol concentration as determined by turbidimetry and colorimetry, according to ShaF ( I 5 ) , was 413 p.p.m. It is possible that the voltammetric wave resulted from the simultaneous oxidation of phenols and aniline bases, as Rogers and Lord (IS) have s h o m that aniline can be oxidized at a microanode.
Table 11. Determination of Phenol in Ammonia Still Waste Liquor" Other materials resent. CaC12, Fe(CN)s-' S C N - t a r bases, unsaturated hydrocarbons, anzorganic molecules of unknhwn strActure. Phenols consist of phenol, 0 - , m-, and p-cresols, and some xylenols. 1 ml. of liquor added to 100 ml. of buffer a t pH 9.0. Concn. in Concn. i n Scale, E I I I I , I d , ED, E M . Buffer, Liquor, No. pa. Volt pa. Volt T70lt .l.Iole/L. Mole/L. P.P.%I. 3 0 485 0 99 0 31 0 60 6 6 X 10-5 6 6 X 10-0 712 64 65 J 0 493 0 85 0 35 0 60 5 7 X 10.6 5 7 X 10-6 620 Av. 666 A46 riiinisheri b y Joceph ?, Shan from Koppers Co., Cobe Plant.
-
__
.~
~~
~
lSssuming Equations 2 descrihe the reaction thehalf-wave potential is given by Equation 4:
+ 0.059 ( log & + log [ K D+ ( H f ) J - log K n ) K
El/* = Eo
(4)
In both cages the equation of the xvave is given by Equation 5:
E
Ei/2 + 0.059 log I / ( l d - I )
(5)
K , is the proportionality constant between diffusion cui rent and concentration and may be evaluated as follows: K1 = ARTAo/FI
(6)
where 11 is the surface area of the electrode, 1 is the theoretical diffusion layer thickness, and the other symbols have their usual meaning.
_._______-
DISCUSSION
Oxidation of Phenol in Terms of Free Radicals. Oxidation of phenol by ozone ( 7 ) and by organic peroxides ( 4 ) gives rise to similar oxidation products. It is likely that the oxidation proceeds through free radical intermediates ( 2 , 16). The fact that electrolysis of phenol produces the same products as oxidation by ozone suggests that electrolytic oxidation proceeds through the same free radical intermediate. This idea is corroborated by the voltammetric proof of one-electron transfer which can result only from the formation of the phenoxide free radical. h certain degree of stability for this free radical should be derived from resonance between the following structures:
04
Figure 3.
7
Tn o possible courses for the reaction have been considered: Either molecular phenol is oxidized directly, as in Equation 1 C B H ~ O ---f H C6H50.
+ H+ + e
(1)
or phenol first ionizes and the phenoxide ion is oxidized as in Equations 2: fast C,H,OH C6H50H+ (2a) slowC~HSOCBH~O e (2b)
e
+ +
ilssuming Equation 1 describes the reaction, the half-wave potential as a function of pH is given by Equation 3 (8):
EI/Z = Eo
+ 0.059
(3)
II
Experimental Values of
El/?
Table 111. Comparison of Voltammetric Data with Apparent Oxidation Potential Values EM. A.O.P., ED, El, 2 , 1
I . The diffusion current is proportional to total phenol concentration and, over the range studied, independent of pH. 2. The half-wave potential's independent of pH a t values where pH is less than pKn. 3. One electron is transferred in the step responqihle for the u ave.
9
7
PH
B 15
The number of products, from ozone oxidation and electrolytic oxidation, where hydroxyl and phenyl groups were substituted in the ring, would indicate a greater contribution of structures with the free el~ctronin the ring. Theoretical Consideration of Wave. Any equation derived for thc wave must account for the following observed facts.
1
5
Volt 0.780 0 460
T'ol t
0.65 0.42
Volt 0.755 0.60
Volt
0.84 0.68
Kz is a constant involving the diffusion coefficient of the phenoxide free radical and the rate constant for the decomposition reaction a t the electrode surface. Yo means for evaluating K Zhas as yet been found. Equations 3 and 4 are not the same, since a t p H values under 8-i.e., where [H+] >> Kn-the values of E differ by 0.059 PKD, or about 0.6 volt. If a means could be found to evaluate K2, it should be easy to distinguish between these two equations. Equation 3 predicts that half-wave potential, El, *,will vary linearly with pH for all values of pH. Equation 4 predicts that the D phenol is 9.94 (S)] is variation is linear until pH = ~ I Z D[ ~ Kfor reached, a t which point Ell2 becomes pH independent. Figure 3 shows a plot of experimental values of Ell2 at various p H values, and a plot of Equation 4. The agreement is good in view of the fact that the over-all reaction is irreversible. This shows that Equations 2 give a more correct representation of the reaction than does Equation 1. Equation 5 is correct, since a plot of log I , ( I d - I ) versus E gives a straight line. K , is valued a t 4 X 10-1 coulomb liter per mole using Equation 5 where A0 = 357 mhos per mole (14) and 1 = 0.25 mm. ( I O ) . The average experimental value of K1 is 1.4 X coulomb liter per mole. This differewe is partly due to the fact that 1 was assumed to be the same for phenol as for ferricyanide ion, and that diffusion was assumed t o be independent of ionic strength up to a value of a t least 0.2. Correlation with Apparent Oxidation Potential. Conant ( 3 )determined the apparent oxidation potential of phenol a t pH values of 5 and 7. As was mentioned earlier, Muller ( I d ) has
ANALYTICAL CHEMISTRY
1358 shown that the apparent reduction potential corresponds to the decomposition potential, ED, on the polarographic wave of a compound. Table I11 shows that a t a p H of 7 the apparent oxidation potential is close to ED. At a p H of 5 the agreement is not so good, but the apparent oxidation potential corresponds to a potential on the v o l t a m e t r i c aave. ANALYTICAL kPPLICATIONS
Qualitative. The half-wave potential of a phenol at a known pH value can be used for qualitative identification. In cases where more than one phenol have the same half-wave potential, further characterization ~ o u l dbe necessary. Qualitative identification of several phenols in a mixture would be possible only if the individual half-wave potentials were a t least 0.2 volt apart. A wave observed to cover a potential range of more than 0.2 volt would indicate a mixture of phenols whose half-wave potentials R ere too close to yield separate waves. Quantitative. In cases where it is necessary t o know the approximate concentration of a phenol, voltammetry would find application. This might be the case with drinking Jvater or industrial waste products, where concentrations of phenols must be held below some maximum value. While the accuracy of this method is only within &14%, it offers a great advantage in its speed and simplicity. -4total of 6 minutee is required to run a sample. Samples containing as little as 0.1 p.p.m. of phenol could be run directly, while more dilute samples could be run if first subjected to processes designed to concentrate the phenol. Antioxidants. Egloff and coworkers (11) have found that effective inhibitors have critical oxidation potentials within the region of 0.6 to 0.9 volt. The half-aave potential of a new
phenol could be determined in much less time than the critical oxidation potential and would show whether more detailed testing for use as an antioxidant was warranted. LITERATURE CITED
Bowman, Ji. H., Ph.D. theais, Vniversitg of Pittsburgh, 1950. Campbell, T., and Coppinger, G., J . ii?n. CRem. SOC.,74, 1469 (1952). Conant, J. R.,and Pratt, AI. F., C‘hem. Recs., 3, 1 (1926). Cosgrove, S. L., and Waters, W. A , , J. Chem. Soc., 1949, 3189. Fieser, L. F., J . A m . C’henr. Soc., 52, 5204 (1930). Fletcher, W. E., Ihid., 68, 2726 (1946). Gibbs, H. D., Philipp. .I. Sci., 4, 133 (1910). Hedenburg, J. F., M.S. t,hesis,t-nirersity of Pitkburgh, 1952. Julian, D. B., and Ruby, W. R., J . A m . Ckeni. Soc., 72, 4719 (1950). Laitinen, H. A , , s n d Rolthoff, I. 3f., J . Phus. Chnn., 40, 1061 11941). Loury, C. D., Jr., Egloff, G., hlorreli, J. C., and Dryer, C. G., Ind. Eng. Chem., 25, 804 (1933). Muller, 0. H., a n d Raumberger, J. P., ,J. A m . Chein. SOC.,61,590 (1939). Rogers, L. B., and h i d , S 8 , Ji., Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, March 5, 1952. Scudder, H., “Electricvtl Conductimty and Ionization Constants of Organic C‘omponrrrls,” p 246, S e \ \ York, D Van Nostrand Co.. 1914.
Shaw, J. A., private coniinunication. Taube, H., J . Am. Chem. Soc., 63, 2453 (1941) for review June 17. 1952. -4ccepted June 26, 1953. Presented at the Pittsburgh Conference on Analytical Chemistry aud Applied Spectroscopy, Pittsburgh, Pa., March 6. 1952. Contribution 862 from the Department of Chemistry, University of Pittfiburgh. Abstracted from the thesis of John F. Hedenburg presented to the Graduate School of the University of Pittsburgh as partial fulfillinent of the requirements for the M.S. degree. RECEIVED
Separation of Phenols by Partition Chromatography T. R . SREENEY AND J. D. HULTRIAN Naval Reseurch Laboratory, Wunhington 25, D. C;
T
H E so-called tar acids coiist,itute B mitjor and, to the woodpreserving industry, a possibly important fraction of high temperature coal tar creosote. In considering the separation and identification of the const,ituents of a certain fraction of t8hetar acids, one that consisted almost entirely of n misture of phenols, the technique of partition chromatography appeared to be applicable. The method of partition chromatography, originally developed by Martin and Syngr ( 2 ) for the separation of t.hc acetyl derivatives of amino acids Tvhich were obtained from protein hydroIj-zates, consists in partitioning a solutr, between two inimisc.it)le solvents, one of which is immoililized by adaorhiiig it on an iiiert solid. In practice, t,he inert solid holding the immobile solvent is packed in a suitable chromatographic tube and, after the solute is placed on the column, the mobile solvent, is percolated through the column, whereupon t,he solut,eis partitioned between the two solvents according t o its partit,ion coefficient.. Under proper conditions the solute will pass through the column in a small zone and eventually emerge from the bottom. When the solute consists of a mixture of compounds, their respective ratcs of passage through the column are a function of their partition coefficients as well as the relative amounts of mobile and immobile phases and hence a separation may often be achieved. For the present invedtigation the use of a system such as thr water (on silicic acid)-cyclohexane system described by Zahner and Swann (6) for the separa.tion of phenol from petroleum cresylic acid seemed particularly applicable, since cresylic acid is a
trade nAmP u v d to tlrqlgnstr a C I ude mixture consisting largely of phenolic ( onipourids but contaminated with neutral oils, sulfur conipounds, and nitrogen compounds. Furthermore, these workers indicated that their system could also be used for the separation of homologs of phenol. Actually, becauw the solubilities of the materials studied appeared to be about the same in iso-octane as in cyclohexane and because pure iso-octane was more readily available, the present study x-as made on the system water (on silicic acid)-iso-octane. In order to employ this system for the separation and identification of phenols in unknown mixtures it was necessary t o study first the behavior of a number of known representative phenols. This paper reports the effect of changes in the r a t e r content of the silicic acid support on the rate of movement of phenolic compounds thiough the column and on their rerolution. EXPERIBZENTA L
The chromatograms in this study were obtained using 12 grams of silicic acid as the support, varying quantities of water as the immobile phase, and isc-octane as the mobile phase. The sample solutions of the individual compounds used contained 1 micromole of solute per ml. (solvent ieo-octane); solutions of mixtures of compounds contained 1 micromole of each per ml. Materials. Phenol, Merck reagent grade. 1-Naphthol, 2-naphthol, p-phenylphenol, o-phenylphenol, and p-tert-butylphenol were recrystallized from water; melting point,s agree with literature. 2,6-Xyle~ol was purified by steam distilling twice: melting point 141.5-143’ C.
