1134
The Journal of Physical Chemistry, Vol. 83, No. 9,
S. Steenken and P. Neta
7979
Electron Transfer Rates and Equilibria between Substituted Phenoxide Ions and Phenoxy1 Radicals' S. Steenken' and P. Neta" Radiation Laboratory, University of Notre Dame, Notre Dame, Indiana 46556 (Received October 19, 1978) Publication costs assisted by the U.S. Department of Energy
The rate constants for electron transfer from a series of substituted isomeric dihydroxy- and diaminobenzenes to different substituted phenoxyl radicals were measured by observing the decay or buildup of one of the radicals involved. In many cases the electron transfer reactions were reversible and the equilibrium constants could be calculated from the individual rate constants for attainment of equilibrium and from the concentrations of the species involved at equilibrium. From the equilibrium constants the one-electron redox potentials for 15 individual Q-./Q2-pairs were determined, using the value for hydroquinone (23 mV at pH 13.5) as a reference. The potential for catechol (43 mV) is near that of hydroquinone, resorcinol is oxidized much less readily (300 mV), while phenol is even a weaker reductant (>500 mV). Methyl, methoxy, and hydroxy substituents decrease the redox potentials while acetyl and carboxyl substituents increase these values. Ascorbate has a potential (15 mV) similar to that of hydroquinone, while TMPD (82 mV) and p-phenylenediamine (183 mV) are less easily oxidized.
Introduction Electron transfer reactions have been extensively studied by spectrophotometric pulse radiolysis. Recently, this technique was further applied to the determination of one-electron redox potential^^-^ from equilibria such as
The equilibrium constant K1 is determined either from the kinetics ( k J 4 - J or from the equilibrium concentrations ([Q1][Qz-]/([Q1-][Qz])). From K1,the difference between the potentials (E1)of the two systems Q1/Q1- and Q2/Q2can be calculated. After establishing several reference potentials, the values of El had been determined for a series of quinones and nitro corn pound^.^-^ From the value of El and the classical two-electron reduction potentials (E,) for Q/H2Q one can calculate the potential E' for the second reduction step (i-e.,for Q-/Q2-). However, experimental evaluation of E2 from equilibria such as Q1-
k + QZ2-2 Q12- + QL k-2
has not been reported. Rate constants k 2 for the oxidation of p-bromophenoxidel and ascorbates ions by the phenoxyl radical were found to be lo8 M-' s-l and the equilibrium was shifted fully in one direction. In the present study a series of substituted phenols and phenylenediamines have been investigated with the aim of establishing the occurrence of equilibrium 2 and measuring the redox potentials for these compounds. Since oxidation of such compounds is an important process in biological systems and is frequently used in industry, the knowledge of the oxidation potentials of some model compounds was considered t o be of interest.
-
Experimental Section Phenols, and especially hydroquinones and catechols, are extremely sensitive to oxygen in alkaline solutions. In order to minimize thermal oxidation, all solutions were prepared in the following manner and were not allowed to come into contact with air. First, the water was bubbled with NzO for -30 min and the proper amounts of phenolic 0022-3654/79/2083-1134$01 .OO/O
solutes were then added. After further bubbling, the appropriate weights of KOH were introduced rapidly. The solution was then flowed through the irradiation cell. The pulse experiments were begun only after thorough flushing of the cell to remove the initial volume of solution. Using this procedure, the solutions remained unchanged for over 1h while the kinetic measurements were usually completed within less than 15 min. The kinetic spectrophotometric experiments were carried out using the computer-controlled pulse radiolysis apparatus described p r e v i o ~ s l y .The ~ ~ ~radiation source in the present study was an ARC0 LP-7 linear accelerator. The pulses used were usually of 54s duration and supplied a dose to produce 2-4 pM of radicals. The water was doubly purified by a Millipore Milli-Q system. The compounds used were of the purest grade commercially available. The inorganic compounds, phenol, and resorcinol were Baker Analyzed reagents; catechol, 4-methoxyphenol, and o-phenylenediamine were obtained from Eastman Organic Chemicals; durohydroquinone and 1,2,4-trihydroxybenzene were from Pfaltz and Bauer; and the rest were from Aldrich.
Method The method used in the present study is basically similar to that described previously in detaiL5 The main difference is that here solutions of the chemically reduced forms of the system are used and a one-electron oxidation has t o be carried out before the establishment of equilibrium 2 can be followed. For this purpose N 2 0 saturated solutions were used so that all the primary radicals of water radiolysis are converted into OH radicals or 0- radicals in strongly alkaline solutions. The latter have been suggested to react by electron transferio C6H60- + 0-+ H@
-P
C6H50 + 20H-
(3)
while OH is known to add to the aromatic ring1' C6H50-
-
+ OH
H0C6H50-
(4)
This addition, however, is followed by acid- and basecatalyzed elimination to produce the phenoxyl radica11i*'2 HOC&o-
-
C6H.50
0 1979 American Chemical Society
OH-
(5)
Equilibria betwjeen SubIrtituted Phenoxide Ions and Phenoxy1 Radicals
As a result, all the primary radicals produce phenoxy1 radicals which can then participate in electron transfer (reaction 2) or decay by radical-radical reactions. However, reactions 3 and 4 ma:y not be quantitative when the phenol contains substituents such as methyl, amino, or methoxyl. In the latter case, demethoxylation was shown13to have varying degrees of participation, and the system cannot be, therefore, used for quantitative evaluation of equilibrium 2. To overcome this difficulty, it is advantageous to use the formylmethyl radical for oxidation by electron transfer without the involvement of H ab~traction.'~The CHzCHO radical is produced from ethylene glycol by HOCH2CH20H+ OH HOCH,CHOH + H 2 0 (6)
---
followed b:y a rapid water elimination in alkaline s01ution'~J~ HOCH,CHOH
OH-
-
HOCH~CHO-
-OH. .
CH,CHO (7)
The resultant radical was found to oxidize various phenols by direct electron transfer14 CH2CH0
+ C6H50- + H+
-
CH3CH0 + C6H56
(8)
By using a large excess of ethylene glycol (-1 M) in alkaline solutions, the direct reaction of OH with the phenol can be prevented and oxidation by reaction 8 predominates. Anilines in general undergo similar reactions to those outlined for phenols and produce the corresponding anilino radicals.16 111the case of. tetramethyl-p-phenylenediamine (TMPD), oxidation by CHzCHO was shown to be rapid14 and was used in the present study for the production of the cation radical. After the initial production of Q1-, or a mixture of Q1and Q2-, these radicals engage in reaction 2. In order to be able to measure equilibrium constants, the attainment of equilibrium has to be fast as compared to decay by radical-radical reactions. In all cases studied, this latter interfering reaction was examined in a one-compound system and its contribution was neglected when its rate was a t least an order of magnitude lower than the rate of achieving equilibrium 2. The kinetics of the approach to equilibrium were determined by following the decay of the Q1- absorption or the buildup of the Q2- absorption. In general, the osemiquinone radicals absorb in the 300-350-nm region, while the meta and para radicals absorb in the 420-450-nm region.13J7 It was, therefore, convenient to study the ortho vs. one of the other isomers but not the meta vs. the para. The calculaition of k z and h-, from the observed kinetics, and the calculation of the equilibrium constant from the concentrations of the two radicals and their parent compounds, were carried out as described p r e v i ~ u s l y . ~
Results and Discussion Preliminary experiments in neutral solutions of catechol and hydroquinone slhowed that the rate of electron transfer involving the neutral compounds is too slow to be observed under pulse radiolysis conditions before Q1- decays. This finding is in fact part of a general trend. It has been shown previously14 that oxidation of phenols and hydroyuinones is much slower than that of their anions. In order to work with the fully dissociated compounds 1:Q2-)most experiments were carried out at pH
The Journal of Physical Chemistry, Vol. 83, No. 9, 1979
1135
13.5. At this pH all the dihydroxybenzenes exist predominantly in the dianion forms:
The values of pKl have been reportedlgZ0for many of the phenols listed in the tables and they are generally in the region of 9.5-11. The values of pK,, however, are known less accurately, but they are in general between 11 and 13. The results summarized in Tab!e I represent in most cases the electron transfer between Q1- and QZ2-as given in eq 2. In certain cases Q t - may be partially protonated even at pH 13.5 (e.g., catechol and durohydroquinone). Several experiments carried out at pH -11 are summarized in Table 11. The agreement between the values in this table and those extrapolated from Table I will be discussed below. The rate constants hz for electron transfer are found to be generally higher when the equilibrium constants K are larger. This trend can be predicted from the Marcus theory as shown, for example, by the quantitative treatment carried out for several quinone/semiquinone systems.21~22 However, a plot of log kz vs. log K from Table I showed a large scatter, which probably results from the wide variety in the chemical nature of the reacting species. Using the equilibrium constant K in Tables I and 11, a set of interrelated redox potentials was calculated. In order to derive the potentials vs. the normal hydrogen electrode ("E) the value for hydroquinone (E2~-lsz= +23 mV at pH 13.5)4was used as a reference. This number has been derived4 from experimental data in neutral and acid solutions and its accuracy depends on the knowledge of pK1 and pK, as mentioned above. The potentials are summarized in Table I11 in the form of reduction potentials of Q- to Q2-. It can be seen from Table I11 that electron-donating substituents enhance the ease of oxidation in the order CH3, OCH3, OH, as expected from their Hammett's substituent constants. The electron-withdrawing group CH3C0 exerts a strong effect in the opposite direction. The effect of the carboxylate group appears to vary with its position on the ring. The potential for resorcinol is quite similar to that of 4-methoxyphenol but much higher than the values for catechol and hydroquinone. This comparison suggests that the radical from resorcinol resembles a substituted phenoxyl rather than a semiquinone radical. A similar conclusion can be drawns from previous measurements of the rates of electron transfer from ascorbate to these radical^.^,^^
The oxidation of p-phenylenediamine (183 mV) is less favorable than that of its tetramethyl derivative (TMPD, 82 mV) and of the hydroquinone dianion. Since the diamines do not undergo acid-base equilibria in the region of pH 7-14, they can be used as references for comparing results at different pH values. The potential for pphenylenediamine determined vs. hydroquinone at pH 13.5 is 183 mV, while that determined at pH 11 (using 57 mV for hydroquinone at this P H ) is ~ 200 mV, in reasonable agreement. The redox potential of ascorbate was measured against that of catechol which was in turn determined vs. hydroquinone both at pH 13.5 and 11. The potentials at pH 11 for ascorbate and catechol are 85 and 139 mV, respectively, as determined from the equilibrium constants and using 57 mV for hydroquinone as a reference. For comparison, the potentials a t this pH can be calculated from the results at pH 13.5 and the pK values for ascorbate and catechol and for their radicals as given in the liter-
1138
The Journal of Physical Chemistry, Vol. 83, No. 9,
1979
S. Stsenken and P. Neta
TABLE I: Electron Transfer Rates and Equilibria at pH 13.5
Q,-+ Q,"
7 k,
Q12-
t
9,-
-2
k,C M-Is - l
Q,"
k,
Kd
catechol 3,4-dihydroxybenzoate 2,3-dihydroxybenzoate pyrogallol p-phenylenediamine resorcinol 4-methoxyphenol ascorbate 2,5-dihydroxyacetophenone hydroquinone 2,5disulfonate
hydroquinone hydroquinone hydroquinone hydroquinone hydroquinone catechol catechol catechol catechol catechol
445e
1000 > 1000
1.9 x 107
1.5 x
-
- l ox5 105
lo6
1.2 x 105
12.7
2.1 35 37 0.3 49 0
- 0.3 28
- 0.7
1.5
- 600 x
105
lo4
- 154
148
900
-1
450 450 565 565 565 565 45 Oe
2.2 x l o y 1.5 X 10' -4.2 x 107
420
7
480 445 565 430
6.6 x lo8 7.5 x 10' 1.5 x 109 -1.0 x 105
x
3.6
>6000 105
3.7 x 105 2.7 x l o 6 -1.2 x 105
6000 56 360
-
-10000 85 1000
405 750
362 964
-
10' 1.6 X
lo6
1.0x 105
a The notations Q, and Q, are used to indicate the irection of the electron transfer as given in the equation The charges in this equation are accurate only for the case where 0th compounds are dihydroxybenzines. In the case of'phenols and phenylenediamines, Q- represents the neutral phenoxy1 or anilino radicals, while for TMPD it represents the cation radical. The wavelengths listed here are those used for the observation of rise of Q2- or decay of Q; and for the determination of the equilibrium concentration. Most of them are the absorption maxima as reported in the literature and verified here except where noted. The experimental error limits in the determination of these rate constants were usually L 10% except where noted. The error limits depended on the individual system and are reflected in the error limits given in Table 111. e Absorption maxima determined in the present study.
TABLE 11: Electron Transfer Rates and Equilibria at p H
- 11" K
k , M-I s-'
Q!
-
phenol catechol catechol p-phenylenediamine p-phenylenediamine o-phenylenediamine
TMPD catechol resorcinol a
Q?
hydroquinone hydroquinone hydroquinone hydroquinone hydroquinone hydroquinone hydroquinone ascorbate
TMPD
PH 11.6 11.0 7.0 11.0
9.0 11.5 11.5 11.1 11.6
k,
k.2
k,/k.,
x io4
> 15
2.2 x 109 -8x lo5
-5
-4x
-1x
500 phenol >500
pH -11 hydroquinone ascorbate catechol p-phenylenediamine resorcinol (pH 13 $6)
t 57d
t85f t139i t 200 i t304i
5 20 13 5
a The potentials are given for reduction of
(t53)e (t98)e ( t1 8 3 4 (t310)e
Q-to Q".
Given vs. NHE using the value for hydroquinone as a reference. T'he error limits reflect only the experimental error limits in the determination of the equilibrium constants K i n Tables I and 11. The accuracy of the absolute values will depend, of course, on the accuracy of the reference number. ' This value differs from that calculated previous1 probably as a result of uncertainties in pK values. lDerived previously4 and used here as a reference value. '? Calculated from the potential at pH 13.5 and the pK values, see text. f Value determined at pH 13.5 and should remain the same at pH 11,see text.
TABLE IV: Calculateda Redox Potentials for Semiquinone/Hydroquinone
ascorbicacid hydroquinone catechol resorcinol
4.5! 9.Ei5 9.45 9.8
11.5 11.4 12.8 11.3
-0.45 4.1 5.0 7.0
0.30 0.459' 0.53 0.72
0.99 1.041' 1.06 1.13
a Using the equation below as giver1 in ref 3. The pK of the radical, from ref 13, 17, and 24. ' From ref 3.
as that determined a t pH 13.5 (TMPD has no acid-base equilibria in this range) a value of 304 mV is calculated for resorcinol. This value is in good agreement with 310 mV calculated from the pH 13.5 result and using the pK values for resorcinollg and its r a d i ~ a 1 . l ~ The same type of calculation can be used to estimate the potentials at pH 7 and 0. The values for hydroquinone have been calculated as 0.459 and 1.041 V, re~pectively.~?~
The Journal of Physical Chemistry, Volt 83, No. 9,
7979 1137
The values calculated here for catechol, resorcinol, and ascorbic acid are given in Table IV. The relative ease of oxidation of the dihydroxybenzenes remains in the same order (para > ortho > meta) as that at high pH and the differences in potential are also somewhat similar in neutral and alkaline solutions. Ascorbate, with a calculated one-electron potential of 0.30 V at pH 7, is a much stronger reducing agent compared to hydroquinone. A t this pH, p-phenylenediamine is also a stronger reductant than hydroquinone, in contrast to the situation a t pH 13.5. In conclusion, it has been demonstrated again that pulse radiolysis can be used for the determination of one-electron redox potentials. In previous studies, the potentials for the first one-electron reduction (E1)have been measured for many quinones and nitro compounds, using these oxidized forms as starting materials. In the present work the reduced forms have been employed. They were oxidized by a one-electron step, followed by electron transfer to reach equilibrium. Redox potentials (E2)were thus determined for a series of dihydroxy and diamino compounds in basic solutions. Direct measurements in neutral solutions could not be carried out because electron transfer reactions involving the neutral phenols are too slow for this purpose. However, the redox potentials of the anions measured in alkaline solutions are useful for estimating the corresponding potentials of the neutral molecules in neutral and acid solutions. The same procedure can be applied for the study of analogous compounds of biological importance.
References and Notes (1) The research described herein was supported by the Office of Basic
Energy Sciences of the Department of Energy. This is Document No. NDRL-1945 from the Notre Dame Radiation Laboratory. (2) On leave from the Institut fur Strahlenchemie, D-4330 Mulheim, Germany. (3) Y. A. Ilan, G. Czapski,and D. Meisel, Biochim. Biophys. Acta, 430, 209 (1976). (4) D. Meisel and G. Czapski, J. f h y s . Chem., 79, 1503 (1975). (5) D. Meisel and P. Neta, J . Am. Chem. Soc., 97, 5198 (1975). (6) P. Wardman and E. D. Clarke, J . Chem. Soc., Faraday Trans. 7, 72, 1377 (1976). (7) R. H. Schuler, P. Neb, H. Zemel, and R. W. Fessenden, J. Am. Chem. Soc., 98, 3825 (1976). (8) R. H. Schuler, Radiat. Res., 69, 417 (1977). (9) L. K. Patterson and J. Lilie, Int. J. Radiat. phys. Chem., 6, 129 (1974). (10) P. Neta and R. H. Schuler, J . Am. Chem. Soc., 97, 912 (1975). (11) E. J. Land and M. Ebert, Trans. Faraday Soc., 63, 1181 (1967). (12) P. O'Neill and S. Steenken, Ber. Bunsenges. fhys. Chem., 81, 550 (1977). (13) S. Steenken and P. O'Neill, J . fhys. Chem., 81, 505 (1977). (14) S. Steenken, J. fhys. Chem., 83, 595 (1979). (15) K. M. Bansal, M. Gratzel, A. Henglein, and E. Janata, J. fhys. Chem., 77, 16 (1973). (16) H. Christensen, Int. J . Radiat. fhys. Chem., 4, 311 (1972). (17) G. E. Adamsand B. D. Michael, Trans. FaraCbySoc.,63, 1171 (1967). (18) A. Albert and E. P. Serjeant, "The Determination of Ionization Constants", Chapman and Hall, London, 1971. (19) G. Kortum, W. Vogel, and K. Andrussow, "Dissociation Constants of Organic Acids in Aqueous Solution", Butterworths, London, 1961. (20) "Stability Constants", Chem. Soc., Spec. Pub/., No. 17 (1964). (21) D. Meisel, Chem. fhys. Lett., 34, 263 (1975). (22) D. Meisel and R. W. Fessenden, J. Am. Chem. Soc., 98, 7505 (1976). (23) P. O'Neill, D. Schulte-Frohlinde,and S. Steenken, J . Chem. Soc., Faraday Discuss.. 63. 141 (1978). (24) G. P. Laroff,R. W.' Fessenden, and R. H. Schuler, J . Am. Chem. Soc., 94, 9062 (1972).
