OXIDATIONOF PEROXIDASE BY ONE-ELECTRON OXIDIZINGAGENTS
Feb. 20, 1956
74 1
ORGANIC A N D BIOLOGICAL CHEMISTRY [CONTRIBUTION FROM
THE
ELDRIDCE REEVESJOHNSON FOUNDATION FOR MEDICAL PHYSICS, UNIVERSITY OF PENNSYLVANIA]
On the Oxidation of Peroxidase by One-electron Oxidizing Agents1 BY R. R. FERGUSSON RECEIVED MARCH 16, 1955 The stoichiometry and kinetics of the action of bromoiridate (a one-electron oxidant) on peroxidase were studied in order The oxidation potential of KJrCle was found as Eo = -0.93 volt, that of KzIrBrR as A? = -0.92 volt. Above pH 7 , hydrolysis occurs. It was found that more than 4 moles of KzIrBra is required to oxidize one mole peroxidase t o complex I. Complex I1 was never observed to be formed from the enzyme directly but always by way of complex I. The results of this study make it seem possible that complex formation occurs by virtue of an intermediate formation of H202, possibly by a Haber-Weiss back reaction or as a product from the reaction of KpIrBra with protein groups such as histidine. The applicability of this hypothesis is discussed. t o shed light on the reaction mechanism.
Increasingly detailed studies of the mechanisms of enzyme reactions have led to many modifications of the general principles to which these systems have formerly been supposed t o adhere. One such principle has been the assumption that hydrogen peroxide is the specific substrate for the iron-porphyrin containing enzymes catalase and peroxidase. The discovery by George2 that the enzymatically active complexes of peroxidase can also be formed when this enzyme reacts with other oxidizing agents has forced a more searching analysis of this assumption. George has shown that peroxidase can be oxidized by a variety of one- and two-equivalent oxidants including K2IrCl6, K & I O ( C N ) ~NaC102, ,~ NaOC1, C102 and O3 to compounds similar to the enzymatically active complexes I and I14J which are, respectively, two and one oxidizing equivalents higher than peroxidase itself. The complexes of peroxidase and catalase may play an important part in the oxidative pathways of living cells. That hydrogen peroxide is produced within the bacterial cell in sufficient quantities to convert catalase completely to its complex I has been recently demonstrated by Chance.6 Nevertheless the inference from George’s work6 that some other oxidizing agent could function in place of hydrogen peroxide needs a critical examination.’ The purpose of this paper is to discuss this suggestion with special reference to the experiments described herein on the oxidation of peroxidase by one-electron oxidizing agents. The following considerations were believed to afford the strongest support for the proposal that (1) This work has been supported in part by the National Science Foundation. (2) P. George, Nature, 169, 612 (1952). (3) This oxidant apparently leads t o the formation of complex I1 only. In addition this compound hydrolyzes very readily in the light. The cyanide ions so formed complex very readily with peroxidase and change its light absorption in those regions of the spectrum where the oxidation process is best studied. The ionization of a spectroscopically observable heme-linked group also affects the absorption in solutions of pH >10 necessary for complex formation with this compound. (4) P. George, Arch. Biochem. Biophys., 46, 21 (1953). (5) P. George, J . Biol. Chcm., 201, 413 (1953). (13) B. Chance, Science. 116, 202 (1952). (7) Recent reviews of the mechanisms of peroxidase and catalase reactions have been provided by P. George and D. H. Irvine, B y i f . J. Radiology, 27, 131 (1954), and B. Chance and R. R. Fergusson.B (8) B. Chance and R. R. Fergusson in “The Mechanism of Enzyme Action,” W. D. McElroy and B. Glass, ed., John Hopkins University Press, Baltimore, Md., 1954, p. 389.
H2OZ is not an intermediary when the aforementioned oxidants react with peroxidase. 1. The reactions between peroxidase and HOCI, HOBr, NaClOz and O3 leading to the formation of the complexes are very fast. If peroxide is produced by a simple non-catalytic mechanism, a bimolecular velocity constant for the initial step of l o 5 mole-’ X set.-' may be required. 2. Peroxidatic activity tests made by measuring the rate of donor (guaiacol) oxidation suggested that peroxidase reacts more rapidly with hypochlorite than with H 2 0 ~ . 5 3 . The rate of formation of the higher oxidation state of metmyoglobin, a compound quite different from peroxidase complex I, is much faster with the one-electron transfer oxidants KzIrCl6 and K3Mo(CN)a than i t is with H ~ O Z . ~ Taking up these points in turn, it suffices to say of the first that Chance’s direct studies’O of the rate of complex I formation using NaC102 as oxidant are consistent with a bimolecular velocity constant no higher than 4 X lo5 mole-’ X set.-' for the initial step; a rate constant of this magnitude would be adequate to ensure that the initial step be rate determining. Chance’s unpublished data suggest that the reaction with NaOC1 is not of the second order. Two-electron oxidative reactions of hypochlorite with velocity constants approaching this order of magnitude are known to occur’’ and a rapid initial reaction which proceeded by oxygen atom or hydride ion transfer might conceivably give rise to peroxide formation. To the second point it can be stated that activity determinations based on donor oxidation can give ambiguous results unless care is taken to ensure no unknown intermediates are rate limiting and corrections are made for the non-enzymatic oxidation of donor. -4 recalculation of George’s work shows that it suffers from some of these faults and in addition does not give a rate of reaction of hypochlorite with peroxidase that is faster than the rate with H202. The following points of George’s paper5 need comment: (a) Of the oxidants listed in Table 11,6only HOC1 shows a rate exceeding that of HzOZ. (9) P. George and D. H. Irvine, Biochcm. J . , 65, xxv (1953), and 68, 188 (1954). (10) B Chance, Arch. Biochcrn. Biophys., 41, 425 (1952). (11) R . E. Connick, TEIS JOURNAL, 69, 1509 (1947).
R. R. FERGUSSON
742
(b) Using the plots of Fig. 7,6 we get initial rates of not more than 0.070 density unit per minute, so that, since
k'
1 1 ___= 10-6 2,6 400 5 10-6 1.1 X 107mole-11. sec.-lat 30" for a0 = 5 X
1 X 60 = ___
x
x
x x
0.77 X lo7 mole-' 1. set.-' a t 30" for a0 = (c) The values that we can compute from George's data (cf. the second do not exceed the velocity constant for HZOZ point above) as obtained from experimental data presented below. George found that guaiacol reacts with HOCl even in the absence of enzyme. Strong evidence for this is also provided by Fig. 7 where the final concentration of tetraguaiacol amounts to only about one tenth of the value calculated for a complete rea.ction of HOC1. (d) The rate of formation of tetraguaiacol is not proportional t o the [HOCl]. A 100% increase in HOCl concentration leads to an increase in rate of a t the most 30y0 (cf. also above under (b)). Chance and MaehlyI2 have recently shown by experiment that donor (guaiacol) oxidation by peroxidase and H202 can be made to depend on kl (formation of complex I ) or k, (reduction of complex I1 by the donor) depending on the conditions of the activity assay
VOl. 78
was prepared by a modification of Theorell's method.I3 The purity was characterized by R.Z. 2.86 and the enzymatic activity by the following values for the rate constants with HZOZand guaiacol, respectively
kl ka
X 107 mole-' 1. set.-' a t 30' 3.1 X lo6 mole-'l. set.-' l a
= 1.03 =
The KzIrCls and KzIrBr6 were prepared by standard methods and estimated by titration against FeSOa.(NHa)2S04. 6HzO. Solutions were made up in 0.01 molar HClO, to prevent hydrolysis. Procedures.-Polarographic measurements were made with a rotating (1800 r.p.m.) gold electrode chosen in preference to platinum because of its higher anodic decomposition potentials. A silver/silver chloride electrode was the reference standard. Spectroscopic measurements were made on a sensitive spectrophotometer equipped to follow fast reactions." Figure 4 was obtained by a sensitive recording spectrophotometer designed by Yang.15 All rate measurements were made a t 25'. Chance's valuesT6for the extinction coefficients of HRP, complexes I and I1 were used.'' The formation of complex I was measured a t 403 mfi. Corrections were made for the absorption of the oxidants a t this wave length. The disappearance of KJrBrs was followed a t 510 mp.
Results Polarographic Measurements.-The oxidation potentials of KdrCle and KaIrBre were measured polarographically in solutions of different pH. There was no evidence for any PH dependence and the potentials were reversible. Table I gives the where results for the chloride. The bromide showed the same constancy of Eo with pH and the measured e = enzyme concn. a0 = initial concn. of donor values of the oxidation potentials in 1 molar HC104 xo = initial concn. of substrate of K21rCle and KzIrBr6 were Eo = -0.93 volt and if klao >> klxo, k l becomes rate limiting; if klxo >> k4ao, Eo = -0.92 volt, respectively. The polarographic k, is rate limiting. Actual experiments a t pH 7 gave the current-voltage curves are given in Figs. 1 and 2. values k l = 0.89 X l o 7 a t 20" and kl = 1.03 X l o 7 a t 30". These values lie within 0.1 volt of the values calcuThe experimental conditions of the activity measurements of George do not allow to decide which of the two rate con- lated from e.m.f. measurements by Dwyer, McKenzie and Nyholm.'8 In solutions more alkaline than stants are measured, since k4no '" klxo. (a0 = 6.7 X kr = 3 X lo5, k4ao = 2 X lo3; kl = 1.0 X lo7,x o = klxo pH 7, hydrolysis of the oxidants occurred as the = 1.0 X lo3). Chance has shown10 that the sequence of curves of Fig. 2 show. The rate of hydrolysis incomplex formation of peroxidase by hypochlorite is very similar to that by HzOz except that the formation of com- creased with the pH of the solution. There are no plex I requires considerably more oxidant. Accordingly no indications that the hydrolyzed species were reduciconclusion can be drawn about the presence of HzO2 from ble a t the electrode. The half-life of KzIrC16in 0.01 such experiments. M borate buffer was of the order of ten minutes. Concerning point three, the higher rate of oxida- K21rBresolutions were found to be more resistant to tion of metmyoglobin by KzIrC16 and K,Mo(CN)s hydrolysis. Because of this and because of its more in solutions of pH greater than 6.8 and 11.O, respecTABLE I t i ~ e l y as , ~ compared with the corresponding rates THE OXIDATION POTENTIAL OF KzIrCle I N SOLUTIONS O F with Hz02 is perhaps the strongest piece of evidence DIFFERENTPH for the non-involvement of H202 in complex forPH Eo (volts) Solution mation, and a one-step electron transfer mechanism -0 -0. 935 1.0 M HClO4 seems plausible with this hemoprotein. But with 0.01 M sodium borate 9.18 - .93 peroxidase and KzIrC16 there is evidence that com5.0 - .94 0.1 M acetate buffer plex I, having two oxidizing equivalents, precedes 1 M potassium chloride -7 - .93 the formation of complex 11. Moreover, the rates of complex formation by HzOa and one-electron oxi(13) A. C. Maehly, in "Methods of Enzymology," Vol. 11, S. P. dants have not yet been compared in the case of Colowick, N. 0. Kaplan, Ed., Academic Press, Inc., New York, N. Y., peroxidase. The results of such studies together p. 801. (14) B. Chance, Rcw. Sci. Instru.. 22, 619 (1951). with titration data and polarographic measure(15) C.C. Yang and V. Legallais, ibid., 26, 801 (1954); C. C. Yang, ments of the oxidants used are here presented. i b i d . , 26, 807 (1954). Experimental Materials.-The horseradish peroxidase (HRP) used in this work was generously provided by A . C. Maehly. It (12) B. Chance and A. C. Maehly, in "Methods in Enzymology," Vol. 11, S. P . Colowick and N. 0.Kaplan, Ed., Academic Press, Inc., in print. New York, N. Y.,
(16) B . Chance, Arch. Biociiem. B i o p h y s . , 41, 404 (1952). (17) Recent work of C. Cecil and A. G. Ogston (Biockcm.J . , 4 9 , 105 1961)) and of H.Theorell and A. Ehrenberg (d. A. C. Maehly, ref. 13) has shown that the molecular weight of H R P is close t o 40.000. T h i s leads to an increase of the extinction coefficients by a factor of 1.1. This change, however, affects the data of this paper very little. (18) F. P. Dwyer, H. A. McKenzie and R. S. Nyholm, J. Proc. R o y . Sac. N . S. W..84, 195 (1951).
Feb. 20, 1956
favorable extinction coeEcients, KzIrBra was the preferred oxidant for most of this work.
-I.oJ'.'.
1000
743
OXIDATION OF PEROXIDASE BY ONE-ELECTRON OXIDrZING AGENTS
.
800
.
.
600
.
.
400
.
. .
200
{,,I,
.
0
Voltage (2000 units E -2.170 volts).
Fig. 1.-Plot of the current us. voltage curve for the polarographic reduction of KzIrCls in 1.0 mole 1.-l HClOd (PH -0): 0 , values of the experiment; 0 , blank values for the supporting electrolyte.
-. 12.0. 4!.
s
t1.1 t0.9 t0.7 t0.5 t03 t0.1 0-0.1 -0.3
-
Voltage (vs AgIAgCI).
10.0c
.g, 8.0. -0 . F
0
6.0-
.
P 4.0-
5- 2.0-
0
E u
E1/2=+0.69volts(VS. AgIAgCI) Eo =-0.92 volts
td
O 1000 /y 000 600 400 200 Voltage (2000 units~-2.102volts).
Fig. 3.-Plot of the current vs. voltage curves KIIrCla a t PH 9.2 (0.01 mole 1.-' sodium borate) to show the effect of hydrolysis. Curves I to IV were measured successively at 10-minute intervals.
disappearance of oxidant increases on increasing the enzyme concentration; again there is no simple relationship apparent. 0
Fig. 2.-Plot of the current os. voltage curve for the polarographic reduction of K11rBrs in 1.0 mole 1.-' HClO, (PH -0).
Formation of Complex 1.-It was verified that in the reaction between KzIrClaor KJrBra and H R P no complex I1 was produced by a bimolecular electron-transfer reaction. At all concentrations used, and for solutions more alkaline than the limiting p H of 6.5, complex I1 was preceded by complex I. I n addition there was a one-to-one stoichiometric sequential transition H R P + complex I + complex I1 + HRP. This is shown clearly by Fig. 4 in which the isosbestic points for these transitions are seen to be essentially time invariant. Kinetics of Complex I Formation and Oxidant Disappearance.-The results can be presented as follows : 1. The disappearance of oxidant follows first-order kinetics. At fixed PH and enzyme concentration, the first-order rate constant does not change appreciably over a range of concentrations of oxidant (see Table 11). 2. The rate of formation of complex I is first order and the measured rate constants a t fixed p H and enzyme concentration are independent of the concentration of oxidant over a wide range (see Table 11). 3 . The rates of both complex formation and oxidant disappearance increase on increasing the PH but there is no simple relationship apparent. 4. The rate of
TABLE I1 RATEOF FORMATION OF COMPLEX I AT DIFFERENTINITIAL CONCENTRATIONS O F KnIrBr6 Initial concn. of KdrBre, 10-8 mole 1. - 1
p H 9.48, borate buffer 0.01 mole HRP, 1.25 X mole 1. -1
39 39 31 23
15.5 fiH 9.18, borate buffer 0.01 mole ].-I, HRP, 1.0 X 10-6 mole 1. -1
15.5 14 20 28
First-order rate constant, sec. -1
0.22 .22 .24 .25 .21 .24 0.076 ,067 .077
RATEOF DISAPPEARANCE O F KnIrBra AT DIFFERENT INITIAL CONCESTRATIONS OF KZIrBrG pH 9.18, borate buffer 0.01 mole 14 0.041 l.-l, HRP, 1.0 X 10-6 mole 17 .037 1. -1 17 .037 20.5 .039 .CIA1 20.5 24 .034 28 ,030
Titration Experiments.-The principal results are presented in Figs. 5-8. These measurements were all made a t PH 9.2 in 0.01 molar borate buffer and they show that: 1. The amount of complex I formed is directly proportional to the amount of peroxidase present initially (see Fig. 5 ) . 2. The amount of complex formed is not directly propor-
P. J. MANNOAND W. H.
808
Vol. 70
JOHNSTON
0 20
Fig. 2.-Change
I
I
~~~~
I
1 2 3 4 Series of washes. Fig. 1.-Relative specific activity (c. min.-l ml.-l X 5) as a function of series of washes for various contact times, where A, B, C, D are 21.5, 43.5, 67.5 and 89.5 hr., respectively. 131-benzene. The amount of iodine-131 was varied to investigate radiation catalysis.
Results and Discussion Exchange Reaction.-A typical exchange study was run a t a concentration of aqueous potassium iodide of F at room temperature. The dependence of the specific activities of iodobenzene of Fig. 1 upon the time of contact is shown in. Fig. 2. For an extremely slow isotopic exchange reaction the usual first-order rate of approach to isotopic equilibrium6 can be expected t o be indistinguishable from a zero-order reaction for the time scale of the laboratory measurement. Thus the rate of isotopic change will be constant giving a linear relationship between the specific activity of a reactant and time from initial non-isotopic equilibrium. The expected linear relation is shown in Fig. 2. The apparent zero-time exchange of Fig. 2 can be explained by a rapid transfer of radioiodine by reaction or exchange with a trace impurity in the iodobenzene. In Fig. 2 the extrapolation to zero time by the method of least squares corresponds to 1.63 X 10-370apparent exchange. This could be caused by as little as 2.0 X lo-' mole fraction of appropriate impurity. When the iodobenzene was pretreated by shaking with a solution of inactive potassium iodide, the apparent zero-time exchange decreased. This indicated that some of the transfer of radioiodine was by chemical reaction as opposed to simple isotopic exchange with an impurity. It was important to find out if the slow transfer of radioactivity to the iodobenzene was caused by (6) (a) H. A. C . McKay, A7alsve, 142, 997 (1938); (b) R . B. Duffield and M. Calvin, THISJOURNAL, 68, 557 (1946); (c) H. T. Norris, J . P h y s . Colloid Clzem., 64,777 (1950); (d) G. M.Harris, Trans. F a r a d a y Soc., 47, 716 (1951).
40 60 80 Time, hr. in the specific activity of the CGHJ with contact time.
a chemical reaction as opposed t o the expected isotopic exchange reaction. In addition to the careful wash-out procedure, 20 ml. of the activated iodobenzene was fractionally crystallized as shown in Table I. I t is apparent that the specific activities of the crystals and supernatants are constant with a satisfactory material balance. A similar check by fractional distillation also showed constant specific activity within experimental error. TABLE I FRACTIONAL CRYSTALLIZATION O F THE TRANSFERRED TO IODOBENZENE CHEMICAL FORMOF
INVESTIGATION BY
Fraction, ml.
Spec. act. (c./m./ml.)
Total act. (c./m.)
0
20.0
2512 2500
Xv. 50,120
s-1 c-1
10.4 9.6
2491 2395
25,906 22,992
s-2 c-2
6.0 3.6
2495 2459
14,970 8,852
s-3
1.3 2.3
2493 2466
3,241 5,672
Fractions
70 Act. balance
97.6
101.0
100.0
c-3
xv. 0, original; S, supernatant; C, crystals.
__
99.5
The question of radiation catalysis in this reaction was investigated by varying the amount of iodine-131 per series by a factor of almost 14. The results are shown in Table I1 in which, for each TABLE I1 INVESTIGATION OF RADIATION CATALYSTS OF THE EXCHANGE REACTION Iia1,a
millicuries
KI, F
Isotopic exchange constant6
Rate of exchangec
Ratio of ratesd
0.15 10-4 28.8 2.88 1.5 0.41 10-2 0.645 6.45 1.2 1.14 10-2 0.511 5.11 1.o 1.58 10-4 18.9 1.89 1. o 2.05 10-4 26.7 2.67 1.4 a At the time of mixing. Units of hr.? X 106. c Units of moles liter-' h r . + X lo9; last digit not significant. Observed to expected (Table 111).
Feb. 20, 1956
O X I D A T I O N O F P E R O X I D A S E BY O N E - E L E C T R O N O X I D I Z I N G
must of necessity be produced if a peroxide is an intermediary in the formation of complex I from a one-equivalent oxidant and peroxidase.22 The Conversion of Complex I1 to Complex 1.When complex I is formed by reaction between KzIrBr6and H R P and reduced to complex I1 by either endogenous or added donor, then a further addition of K2IrBro will reconvert complex 11, apparently directly, back to complex I . Non-specific Protein Oxidation.-The amino acids tyrosine, tryptophan and histidine have been found to react with K21rB1-6 and the rates of is increased. these reactions increase as the ~ J H A t p H 9.2 tyrosine reacts quite rapidly with K21r331-6 (“bimolecular rate constant,” lo3 mole-‘ 1. sec-1). If tyrosine is added to the peroxidase solution the rate of complex I formation by the action of K2IrBr6 is decreased. Thus reactions between protein end groups consisting of these amino acids and the oxidant will account for the observed high ratio of oxidant consumed/complex I formed.
AGENTS
745
kinetic and stoichiometric behavior makes it, unfortunately, impossible to be definite about the nature of the intermediate leading to the formation of complex I. It has frequently been suggested that powerful one-electron oxidants may lead to the formation of hydrogen peroxide by the Hgber-Weiss back reaction, such as Ce+40H- f OH-
+Cef3 + H2Oz
(3)
discussed by Evans and Uri.25 A reaction of this general type involving KzIrBrs and its hydrolyzed species might be expected to proceed quite rapidly because of the favorable energetics involved.26 An alternative theory is that protein end groups such as histidine27 react with the oxidizing agent (KzIrBr6)and form an intermediate that reacts with the iron atom. It should be recognized that these possibilities are a t the present time highly speculative. It is useful to contrast the reactions of peroxidase or ferrimyoglobin with hydrogen peroxide or with one-electron transfer agents such as chloro- or Discussion bromoiridates. Hydrogen peroxide reacts with Although it is very probable that the halogen peroxidase in one of the most carefully studied atoms of K21rBr6and KzIrC16 are inert to substitu- second-order reactions of enzyme and substrate; tion in acidic solution,z3~24 it is clear from the polaro- the velocity constant is very large (lo7 mole-’ 1. graphic measurement presented that hydrolysis sec. -I). With ferrimyoglobin, the reaction proceeds takes place in alkaline media. For this reason the a t less than l / ~ , of o ~this ~ rate and no data have been oxidation potentials of the unhydrolyzed com- published to prove that the reaction is of the secpounds are no longer a very useful guide to the ond-order.lO This paper shows that the reaction of mechanism of the oxidation process; and the polaro- bromoiridate with peroxidase is not second order graphic measurements show no indication of the and is slow (the first-order velocity constant divided lo3 mole-l 1. oxidation potentials for the hydrolyzed species. It by the bromoiridate concentration is worth noting that the effect of p H on hydrolysis set.-'), whereas George and Irvine find that ferriparallels the rate of complex formation. For these myoglobin reacts rapidly with chloroiridate (>lo6 reasons one should be cautious in drawing conclu- mole-’ 1. sec.-l). The nature of the reaction prodsions as to the oxidation-reduction properties of the ucts differs; peroxidase gives the very reactive complexes from the standard oxidation potentials complex I, ferrimyoglobin gives the relatively unreactive compound which is similar to peroxidase comof the oxidants (4.ref. 10). Clear evidence that a direct electron transfer re- plex 111. I n view of these contrasts, the conclusion action does not lead to complex formation is pro- of George that ferrimyoglobin reacts with chloroirivided by the constancy of the first-order rate con- date in a one-electron transfer reaction does not stants for both complex I formation and also oxi- seem to be relevant to the mechanism of the peroxidant disappearance. 9 n d titration data show that dase reaction. more than twice the theoretical amount of K21rBr6 Acknowledgments.-The author would like to is required to form complex I. Thus neither the express his appreciation for the help and encouragetitration nor the kinetic data confirm George‘s sug- ment given him by Professor Britton Chance, Dr. gestion that HzO2 (or H02-) is not a component part A. C. Maehly and the other members of the Johnof the structure of complex I and is not necessary son Foundation. for its formation The extremely complicated
