The treatment of citrus oils with thiourea reduced the terpene content considerably. A higher degree of reduction might be possible, depending upon the ratio of the reactants used. The aging characteristics of the virtually terpeneless oils were compared with the crude citrus oils by exposure to ultraviolet light a t room temperature. After 3 days, the crude citrus oils began to form the usual dark brown residue of the polymerized terpenes. T h e treated citrus oils showed no signs of residue formation even after 7 days of exposure to the test conditions. Laboratory tests indicate that removal of polymerizable compounds helps to protect the oils against flavor deterioration without changing their aroma. Though this would enable
them to meet the specifications for deterpenated orange and lemon oils, we have done no work to evaluate their use in foods. T h e absence of thiourea from the treated oils makes their use possible from a toxicity standpoint. This study was based solely on orange and lemon oils; no conclusions can be drawn for other oils such as peppermint. literature Cited
Miliotis, J. A , , Galinos, A. G., Chim. Chronika 24, 152-4(1959). Swern, David, Znd. Eng. Chem. 47, 216 (1955). RECEIVED for review November 28, 1966 ACCEPTED May 11, 1967
T H E STABILITY OF SEVERAL IRON(III)PHENOL COMPLEXES WALTER R . M A Y , W. R I C H A R D M A T T H E W , AND LEWIS BSHARAH Corporate Research Laboratories, Petrolite Corp., St. Louis, Mo. 631 19
The equilibrium constants for iron(ll1) complexes with phenol and 10 phenol derivatives have been measured in 50 and 75% dioxane-water solutions at pH 2.0 and 25” C. Several of the phenol derivatives which have bulky substituents ortho to the phenol group are rubber antioxidants. It was concluded that these compounds are poor complexing agents and that metal deactivation, therefore, does not enter into the antioxidation mechanism of these types of compounds.
s
or “hindered” phenols are well known as nonstaining rubber antioxidants (70, 14, 19). (These are representative examples of the literature on phenol-type oxidations. Numerous others may be found in the literature and patents.) These compounds usually have methyl, tertbutyl, or phenol groups ortho to the hydroxyl group which create a considerable amount of steric hindrance about the hydroxyl group. One aspect of the general mechanism of antioxidant activity concerns the deactivation of metals which catalyze oxidation (73). This may occur via stabilization of oxidation states of the metal or by strong complexation of the metal ion to its maximum coordination number, thereby preventing further coordination with the hydroperoxide (79). The purpose of the present work was to find if hindered phenols can form iron(II1) complexes which would prevent the iron (111) ions from participating in the oxidation mechanism. Milburn and coworkers (72, 76) evaluated the stability constants of several monosubstituted phenol-iron(II1) complexes spectrophotometrically and potentiometrically a t constant ionic strength and with varying pH. Milburn extended his work to include a n evaluation of the enthalpies and entropies of the iron(II1)-phenolate association (77). Ernst and coworkers (4-7) have also evaluated the stability constants for several substituted phenol-iron(II1) complexes. Milburn (76) has adequately summarized the literature prior to 1955. We have evaluated the stability constants of phenol and 10 substituted phenol complexes with iron(II1) in 50 and 75% dioxane-water solutions. The method developed by Graddon and \Vatton ( 9 ) with a few modifications described below was used. UBSTITUTED
Experimental Materials. T h e water was purified by distilling deionized water. T h e dioxane was Fischer reagent grade material. T h e phenol was Mallinckrodt analytical reagent grade, the p-cresol was purified by distillation, and the p-tert-butylphenol was commercial material purified by recrystallization. T h e 2,2 ’-methylene bis(6-tert-butyl-p-cresol) was obtained from the American Cyanamid Co. The other phenol derivatives were prepared in this laboratory and purified by recrystallization. Reagent grade perchloric acid and sodium hydroxide were used to adjust the pH. Aged iron perchlorate solutions were prepared as recommended by Milburn (76). Spectrophotometric Determination of Equilibrium Constants. T h e technique used to measure the stability constants is based on the method used by Graddon and Watton ( 9 ) to evaluate adducts of pyridine and copper(I1) P-diketone chelates. Absorption spectra and equilibrium constants were determined on a Beckman DK spectrophotometer thermostated a t 25’ C. The absorptions were measured a t the maximum, which occurred around 550 mp on solutions conM total iron(II1) and phenol contents varying taining between 10 and 3000 times the iron content, depending on the solubility of the phenol. The p H was held a t 2.0. The equilibrium constants were calculated in the manner described by Graddon and \Vatton, with the following exceptions. At the concentrations of phenol required to swamp the iron and force it all into the complex form, the solubility of the phenol is exceeded. This condition would be a t infinite ligand to iron(II1) ratio. Therefore, a mathematical approach was required. For two solutions of different phenoliron(II1) ratios and absorptions, the equilibrium constant will not vary. The equilibrium constant expressions for two ligand to iron ratios were equated, and the extinction coefficient for the ratio a t infinity was calculated via a quadratic equation. For each phenol derivative, all possible combinations of data for the extinction coefficient a t the infinity ratio were calculated and the average was used. VOL. 6
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Two other modifications of the procedure were also required to make the method applicable to our problem. T o relate the phenol and iron(II1) originally added to the solution to the phenate ion and iron(II1) ion concentrations available for complexation, the iron hydrolysis and phenol acid equilibria must be considered. Taking Milburn and Vosburgh's value (78) of 3.83 X for the hydrolysis constant of iron(III), 72.2y0 of the iron is calculated to be in the unhydrolyzed form a t p H 2.0. The phenate ion available for complexation is controlled by the acid equilibrium. The phenate ion concentration is equal to K , X [phenol]/[H+] where K, is the acid equilibrium constant. The acid equilibrium constants for the phenols in 50 and 75% dioxane solutions were determined by the method of Bordwell and Cooper ( 2 ) . Results and Discussion
The acidity constant and stability constant data are summarized in Table I. T h e stability constants for the five systems evaluated in the 50% dioxane-water solution were used to evaluate the effect of the phenol basicity on the complex stability in this solvent system. Figure 1 demonstrates the good correlation found between the acidity constant of the phenols (pK,) and the complex stability constant (log K ) . A good correlation was also found between the stability constants and sigma constants compiled by Jaffe (77). A Hammett plot of the log K values us. the sigma constants gave a negative slope, indicating that strongly electron-withdrawing groups reduce the basicity of the phenate ion and the stability of the complexes. These data indicate that the stabilities of the substituted phenol complexes relative to the phenol complex itself are dependent on the electron-withdrawing or releasing power of the substituent on the phenol. Several authors have interpreted similar data to indicate that dative ?r-bonding of the metal [iron(III) in this case] with the ligand takes place (7, 3, 6 , 8, 15), although this has been seriously questioned by other authors (20'). However, we feel that these data are best interpreted on the basis of substituent group electronic considerations alone. The acidity and stability constants from the present work are compared with literature values in Table 11. These data show that our results follow the same trends that are observed in water solution, although in the less polar dioxanewater solution the constants are displaced somewhat. The slope of the log K us. pK, plot for Milburn's data (72), obtained
Table 1.
Stability and Acidity Constants
Acidity Constant, Stability Constant Substituent PK= K log K 50YGDioxane-50yo Water Phenol 11.33 5.82 X 108 9.77 9.07 o-Chlorophenol 10.45 1.18 X 108 p-Chlorophenol 10.87 2.25 X 108 9.35 p-tert-Butylphenol 11.87 3.31 X 10'0 10.52 p-Cresol 11.86 1.80 X 10'0 10.26 757, Dioxane-25yo Water p-tert-Butylphenol 12.57 1.42 X p-tert-Butylphenol dimer 11.07 5.07 X fi-tert-Butylphenoltrimer 10.01 1.57 X >-tert-Butylphenol tetramer 9.67 1 . 3 3 X Di(2-hydroxy-5-octyl) phenyl sulfide 11.07 2.90 X
1010 100 IOIO
10.15 9.70 9.20 10.12
10"
11.46
108
2,2 '-Methylenebis-6-tert-butyl-p-
cresol 4-tert-Butyl-2,2'-bis(3,5-di-tertbutyl-2-hydroxy phenyl)-2,6xylenol
186
11 .64
No complexation
10.10
No complexation
l&EC PRODUCT RESEARCH A N D DEVELOPMENT
t 0 10.00 .
0
1050 o
l
11.00 f "
f 11.10 '
f
12.00 f ~
1210 ' ~
P Ka Figure 1.
Correlation of log K with pK.
Table II. Comparison of Our Data with literature Values Substituent Solvent pKa log K Ref, This work 11.33 9.77 Phenol 50% Dioxane Water 9.98 Phenol 8.20 (76) 9 , 3 5 This work 50% Dioxane IO. 87 p-Chlorophenol 9.42 p-Chlorophenol Water 7.95 (72) 507, Dioxane 11.86 10.26 This work p-Cresol Water 10.25 p-Cresol 9.25 (76) 9.07 This work o-Chlorophenol 50% Dioxane 10.45 Water 8.33 o-Chlorophenol 7.26 (72)
in water, was 1.055 compared with 1.043 for our data. Solutions of 75% dioxane-25% water were used for the phenoltype antioxidants because of solubility problems in the more polar solvents. The p-tert-butylphenol is a slightly stronger base in the 75% dioxane-water solution, although the iron (111) complex is a little less stable. T h e phenol-type antioxidants generally fall into two classes: those in which the phenol has methyl and/or tert-butyl group substituents in the ortho and para positions and those in which two substituted phenols are linked together by a methylene group a t the ortho positions. The phenols chosen for examination in this work were used to test both the effects of substituents in positions ortho to the hydroxyl group and the linking together of several phenols through methylene bridges. A series of p-tert-butylphenols was chosen which contained two, three, and four phenol rings linked together through methylene groups a t the positions ortho to the hydroxyl group. T h e structures of these compounds and their generic names are given in Figure 2. They are referred to hereafter as the p tert-butylphenol dimer, trimer, and tetramer. T h e stability constants for the dimer and trimer are less than for the p-tertbutylphenol, although the dimer and trimer have more oxygens available for complexation. The stability constant for the tetramer is about equal to that of the single phenol. These data are best rationalized by considering steric interferences. The phenol methylene-bridged in the ortho position is probably as effective in preventing complexation by blocking the approach of the iron(II1) ion to the phenolate oxygen as a tert-butyl group. I n the case of the tetramer, the number of
NAME
STRUCTURE
H
8
H D ~ ( ; ~ - H Y D R O X Y - ~ - O C T Y L ) PHENYLSULFIDE
H
0
asD
rings were bridged via a sulfide link, di(2-hydroxy-5-octyl) phenyl sulfide. This compound is theoretically capable of forming a tridentate chelate ring through the formation of two five-membered rings. T h e stability constant for the complex formed with this compound was greater than the constant for the p-tert-butylphenol dimer complex by a factor of 20. T h e enhanced stability is probably due to the presence of the sulfur bridge link which makes possible the formation of tridentate ligand. I n conclusion, it is apparent that hindered phenol-type antioxidants cannot function as complexing agents. Therefore, deactivation of metals which catalyze oxidation is eliminated as a possible mechanism for the antioxidant activity of these materials. Acknowledgment
T h e authors are indebted t o D. B. Merrifield, director of this laboratory, for his generous support of this work. Thanks are also due to C. D. Gutsche, of Washington University, for several fruitful discussions of this work. literature Cited
(1) Basolo, F., Murman, R. K., J. Am. Chem. Sac. 77, 3484 (1955). (2) Bordwell, F. G., Cooper, G. D., Zbid., 74,1058 (1952). (3) Da Silva, J. J. R. F., Calado, J. B., J. Znorg. Nucl. Chem. 28, 125 (1966). (4) Ernst, Z . L., Herring, F. G., Trans. Faraday Sac. 60, 1053 (1 964). ,-I .
sites available for complexation outweighs the steric effects, and the degree of complexation increases. T h e likelihood of bidentate bond formation with the iron(II1) and the dimer, trimer, or tetramer is remote, since a n eight-membered ring would be required. A pair of even more sterically hindered antioxidants, 2,2'methylenebis(6-tert-butyl-p-cresol)and 4-tert-butyl-2,2 '-bis(3,5di-tert-butyl-2-hydroxylphenyl)-2,6-xylenol, were examined. These compounds are analogous t o the dimer and trimer, except that the hydroxyl groups are completely surrounded by bulky substituent groups. I n the case of these compounds, no change in absorption was detected in the solutions, indicating that no complexation had occurred. Molecular models indicate a high degree of steric hindrance which prevents the iron(II1) ion from approaching closely enough to the oxygen to form a bond. A phenol derivative was also examined in which two phenol
( 5 ) Zbid., 61, 454 (1965). (6) Ernst, Z . L., Menashi, J., Zbid., 59, 1794 (1963). (7) Zbid., p. 2838. (8) Falk, J. E., Phillips, J. N., Nature 212, 1531 (1966). (9) Graddon, D. P., Watton, E. C., J. Znorg. Nucl. Chem. 2 , 49 (1961). (10) Ingold, K. U., Chem. Reu. 61, 563 (1961). (11) Jaffe, H. H., Zbid., 63, 191 (1963). (12) Jabalpurwala, K. E., Milburn, R. M., J . A m . Chem. sac. 88. 3224 (1966). (13) 'Lee, L.-H., 'Stacy, C. L., Engel, R. D., J. Appl. Polymer Sci. 10, 1699, 1717 (1966) (extensive review articles cited). (14) Lloyd, W. G., Zimmerman, R. G., Dietzler, A. J., IND. ENG.CHEM.PROD.RES.DEVELOPMENT 5,326 (1966). (15) May, W. R., Jones, M. M., J. Znorg. Nucl. Chem. 24, 511 (1962). (16) Milburn, R. M., J. Am. Chem. Sac. 77,2064 (1955). (17) Zbid., 89, 54 (1967). (18) Milburn, R. M., Vosburgh, W. C., Zbid., 77, 1352 (1955). (19) Scott, G.,Chem. Znd. 1963, p. 271. (20) Yingst, A,, McDaniel, D. H., J. Znorg. Nucl. Chem. 28, 2919 (1966). RECEIVED for review March 20, 1967 ACCEPTED April 24, 1967
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