2459
One-Electron Reduction Potential of Riboflavine
One-Electron Reduction Potential of Riboflavine Studied by Pulse Radiolysis' Dan Meisel" and P. Neta Radiation Research Laboratories and Department of Chemistry, Melion Institute of Science, Carnegie-MellonUniversity, Plttsburgh, Pennsylvania 152 13 (Received July 2 1, 1975) Publication costs assisted by Carnegie-Mellon University and the US. Energy Research and Development Administration
The one-electron reduction potential of riboflavine has been determined at several pH values between 6 and 12 using the pulse radiolysis technique. Solutions containing riboflavine and either duroquinone or 9,10-anthraquinone-2-sulfonicacid were irradiated under reducing conditions to produce their semiquinoidic radicals. Electron transfer from one of the radicals to the parent molecule of the other then took place until equilibrium was achieved. Spectrophotometric measurement of radicai concentrations at equilibrium enabled determination of the equilibrium constant and redox potential. The reduction potential of riboflavine at pH 7 was found to be E7l = -0.292 f 0.005 V. The dependence of the reduction potential of riboflavine on pH was found to be in agreement with the dependence expected from the known pK values for riboflavine and its semireduced form.
Introduction The existence of the semiquinoidic form of riboflavine and other flavines was one of the first examples to establish the theory of Michaelis that reduction of organic compounds takes place by two distinct consecutive one-electron redox reaction^.^,^ Since then extensive studies on the semireduced level of flavines provided information regarding their optical, ESR, and reactivity characteristic^.^-^ Because of the transient nature of these species, the pulse radiolysis technique was found to be very helpful in the determination of their optical absorption spectra and dissociation constants.8 Pulse radiolysis was recently shown also to be a powerful technique by which one-electron redox potentials of short-lived radicals can be directly mea~ured.~JO Thus far the redox potentials for the OdOz- and several quinone1 semiquinone couples as well as the one-electron reduction potentials for some dozen nitroaromatic compounds,ll including several radiosensitizers, have been measured. We feel that this method may be applied to the determination of one-electron redox potentials for more complex compounds such as the components of the biological electron transport system. Riboflavine seems to be a favorable example for demonstrating the capability of this method. Comparison can be made with the results of Michaelis et al.233which were obtained by the indirect method of conventional potentiometric titration, Method The method adopted in this study consists of measurement of the equilibrium constant of the electron transfer reaction between the semiquinone of riboflavine (HZRf) and another reference acceptor (A). H2Rf
+ A a HRf + A- + H+
(1) HzRf is the uncharged form of riboflavine at the semiquinoidic level for which the following acid-base equilibria have to be considered: H&f+ HzRf
PKRI= 2.3
e
~ K R=Z8.3
1J
H&f
+ H+
HRf-+H+
Using the same notation, the following acid-base equilibria apply for the oxidized form
All the pK values are taken from ref 8. Once K1 is determined the potential difference between the two half-cell reactions (reaction 1) can be calculated. If the one-electron redox potential of the acceptor is known the redox potential of the other half-cell can be calculated. Two reference acceptors were used in this study. Duroquinone (DQ, E71 = -0.235 V) was used in the pH range of 5-8, and 9,lO-anthraquinone-2-sulfonic acid (AQS, = -0.380 V) was used at higher pH val~es.~-ll The subscript in the potential notation denotes the pH while the superscript stresses the first electron reduction reaction. All the potentials are given against the normal hydrogen electrode and represent standard potentials in the sense that the ratio of concentrations of the oxidized and the semireduced forms is 1 ([Ox]/ [Sen>]= 1). The radicals were produced by irradiation of aqueous solutions containing 0.1 M sodium formate which serves as a scavenger for H atoms and OH radicals and produces COZby its reaction with these species. lrradiation was carried out using 2.8-MeV electrons from a Van de Graaff accelerator given in pulses of 0.5-1-psec duration and yielding 1-2 pM total radical concentration. At the end of the pulse the only radicals present are the reducing radicals eaq- and COz-. These radicals will reduce riboflavine to give its semiquinone in less than 10 wsec8 under our experimental conditions (0.5-5 X lo-* M riboflavine). The reversibility of the reduction reaction was checked by monitoring the spectra after the decay of the semiquinoidic radicals. These radicals disproportionate to give the oxidized and the totally reduced forms and the final spectrum indeed showed recovery of half of the riboflavine initially consumed. The protonation of the semiquinoidic form took place too rapThe Journal of Physical Chemistry, Vol. 79, No. 23, 1975
2460
Dan Melsel and P. Neta TABLE I: Calculation of Apparent Equilibrium Constant K1' a t pH 7.W
[-I/
[Ribo-
flavine], 111 5 X 5 X lom4 5 x 10'~ 5 x 10'~ 5 x 10-4 2.5 x 10-4 2.5 x IOm4 2.5 x lom4
0
c
e0
T I
:pol .-> 2.5
oH.59
.
?
1.5
1.0
.
.5
.
[DQ], 111 4.0 X 8.3 X 1.3 x 1.7 x 2.5 x 3.2 x 6.6 x 1.3 x
[W-I/
[DQ] [Sem] 12.5 6.0 3.8 2.9 2.0
io-* 10'~ 10-4
10-5
7.8
lom5
3.8 2.1
IOm4
8.3 1.54 2.5 3.4 4.6 1.2 2.4 3.8
Kl'
10.4 9.3 9.9 10.0 9.2 9.4 9.1 8 .O K,' = 9.4 1 E,' = -0.292 V
*
a All solutions contained 0.1 M sodium formate and were purged with Nz. The concentrations of the radicals were derived from measurement of the optical absorption at 570 nm. E$ = -0.235 V for duroquinone was a s ~ u m e d . ~
Approach to equilibrium in the system H2Rf 4- DQ
HRf + DQ- + H+. Deoxygenated solutions containing riboflavine, duro-
Flgure 1.
quinone, and 0.1 M sodium formate were irradiated and the changes in absorbance recorded. (a) [DQ] = 6 X lov5 M, [HRf] = 2.5 X M, pH 6.9 phosphate buffered, absorbance recorded at 445 nm. Bleaching of riboflavine is observed within 5 Wsec after the pulse. The riboflavine chromophore is partially restored while the abM, sorbing durosemiquinone is produced. (b) [DQ] = 2.4 X [HRf] = 5 X M, pH 5.9 phosphate buffered, absorbance recorded at x 570 nm. Partial formation of riboflavine semiquinone is observed immediately after the pulse, followed by slow formation of an additional amount via electron transfer from DQ-. idly to observe, probably due to the high concentration of the buffer M phosphate in the region of pH 6-8). Results and Discussion When both riboflavine and the other acceptor are present in solution they compete for the primary reducing radicals. However, after this reaction is complete, but before the semiquinones decay, another reaction is observed. This behavior is shown in Figure l a where the absorption a t 445 nm, the peak of the durosemiquinone, is followed for 200 p e c after the pulse. The initial bleaching of riboflavine is seen in this figure to be gradually replaced by absorption. We attribute this change to the formation of durosemiquinone by reaction 1. This formation followed a pseudo-firstorder rate law linearly dependent on the concentration of duroquinone. From this dependence k l = 2.5 X lo8 and k-1 = 3.0 X IO7 at pH 6.9 have been determined (cf. ref 11 for details). When a similar experiment was carried out at pH 5.9 reaction 1 proceeded in the opposite direction, i.e., the initial absorption of the riboflavine radical was lower than its equilibrium value and Figure I b shows its relaxation to the equilibrium state. This is to be expected since the pK, of durosemiquinone is and thus its redox potential should not change to any measurable extent in going from pH 6.9 to 5.9 while the reduction potential of riboflavine should change by about 60 mV. We found it more accurate to calculate the equilibrium constant of reaction 1 from the yields rather than from the rate constants. This was accomplished using [OX] [A-] [OX]D - DSem K1'= --(1) [Sem] [A] [A] D - D A The Journal of Physjcal Chemjstry, Vol. 79, No. 23, 1975
O t
2
-a'=-0.2
-
W
-0.4 -
i
-0.6
PH Dependence of the first reduction potential of riboflavine on pH. Experimental points were determined against either DQ (0) or AQS ( 0 )as described in the text. The solid line was calculated using eq Ill and the pKa values for riboflavine and its radical, as reported in ref 8. The pKa values used are indicated by the arrows. Figure 2.
where Ox and Sem are the oxidized and semiquinoidic forms of riboflavine regardless of their protonation state and D is the absorbance at equilibrium. DSemand DA- are the absorbances of the semiquinones of riboflavine and the reference, respectively, obtained after pulsing solutions containing each of them in the absence of the other. K1' rather than K1 was calculated since Ehl is easily calculated from this apparent equilibrium constant. hEhl
=
In K ~ '
(11)
Each of the apparent equilibrium constants at each pH was measured using at least two different concentrations of riboflavine and four of the reference quinone. The yields were measured at two wavelengths whenever possible. Table I presents a typical calculation of K1' and shows that its value is indeed constant over a wide range of concentrations. Figure 2 displays the experimental results of E h l for riboflavine at several pH values. E h l in neutral solution was measured several times under different experimental conditions and the average value obtained is E,1 = -0.292 V.
2461
Temperature-Jump Relaxation Amplitudes Taking this value and the pK,’s for the oxidized and semiquinoidic forms measured previous19 one can calculate the pH dependence of E h l using
ered. Our results favor the value of PKR*= 8.38rather than the one proposed by Michaelis? The results presented in Figure 2 and the close agreement with the potentiometric data give a very high level of confidence in the determination of redox potentials of radicals using the pulse radiolysis technique. We now feel that for systems in which equilibrium 1 can be established the potentials can be measured with an accuracy of f 5 mV, provided that a suitable reference acceptor can be found.
References and Notes
The calculated line is shown in Figure 2 to fit quite well to the experimental results. The measured values for E h l at pH 11 and 11.9 are higher by ~ 2 mV 0 than the calculated values. This shift in potentials may have resulted from the use of a reference system different than that used at the lower pH. The value of E$ = -0.292 V is comparable with the value of -0.285 V obtained originally by Michaelis et aL2 The discrepancy in his other work3 remains unresolved.13 Since in our work equilibrium was measured before any decay took place, by disproportionation or by formation of a “quinhydron type” dimerization, no complications resulting from such a dimerization have to be consid-
(1) Supported in part by the U.S. Energy Research and Development Administration. (2) L. Michaelis, M. P. Schubert. and C. V. Smythe, J. Biol. Chem., 116, 587 ( 1936). (3) L. Michaelis and G. Schwarzenbach, J. Biol. Chem., 123, 527 (1938). (4) H. Beinert, J. Am. Chem. SOC.,78, 5323 (1956). ( 5 ) B. Holmstrom, Photochem. Photobiob, 3, 97 (1964). (6) A. Ehrenberg, F. Muller, and P. Hemmerich, Europ. J Biochem., 2 , 286 (1967). (7) A. Knowles and E. M. F. Roe, Photochem. Photobiol., 7, 421 (1968). (8) E.J. Land and A. J. Swallow, Biochemistry, 6, 21 17 (1969). (9) D. Meisel and G. Czapski, J. Phys. Chem., 70, 1503 (1975). (IO) Y. Nan, G. Czapski, and D. Meisel, Biochim. Biophys. Acta, submitted for publication. (11) D. Meiseland P. Neta, J. Am. Chern. SOC.,97, 5198 (1975). (12) K. B. Patel and R. L. Wilison, J. Chem. SOC.,Faraday Trans. 1, 69, 814 (1973). (13) W. M. Clark, “Oxidation Reduction Potentials of Organic Radicals”, Wllliams and Wilkins, Baltimore, 1960, note f p 445.
Temperature-Jump Relaxation Amplitudes for Single-Step Processes. The Methyl Orange System James Muirhead-Gould and John E. Stuehr. Department of Chemistry, Case Western Reserve University, Cleveland, Ohio 44 106 (Received June 16, 1975) Publication costs assisted by the National Institutes #f Hea/th
Equations are developed for the amplitudes of spectral changes accompanying the concentration perturbations in a temperature-jump relaxation experiment. We show that the conditions for optimizing the photometric signal in terms of stoichiometric concentration and pH are quite different from those for optimization in terms of extinction coefficient and that the conditions in fact depend upon the properties of the system under study. Amplitude data at 2 5 O and I = 0.1 M are presented for the Methyl Orange proton-transfer system as a function of pH and stoichiometric concentration. A computer minimization procedure fitted the data to the two parameters pK and AHo. The resulting values are in agreement with those obtained by conventional thermodynamic measurements. The transient relaxation kinetic techniques were developed, by Eigen and his associates, in the rnid-1950’~.’,~ The theory of relaxation techniques has focused principally on the time course3’*of the perturbations: the calculation of relaxation times for single-step and multistep processes. The calculation of relaxation amplitudes has received little attention. Although they do not provide kinetic information (Le., rate constants), they may be used5 to distinguish among possible mechanisms. The temperature-jump (T-jump) method is the most widely used relaxation technique. In its most common form, Joule heating is used to heat a solution quickly (with-
in -IO+ sec), and thc resulting concentration changes are detected by optical means. Eigen and DeMaeyer6 defined the photometric signal as
s = SIII (1) where 61 is the change in light intensity and I is the transmitted intensity. In practice, this can be misleading since S refers to no well-defined physical quantity. If the transmitted intensity is low (a highly absorbing solution), the signal could be large in terms of eq 1 yet be completely undetectable. The photometric signal for a one-step reaction was shown6 to be The Journal of Physical Chemistry, Vol. 79, No. 23, 1975
