JOURNAL OF COLLOID SCIENCE 11, 327-339 (1956)

A KINETIC STUDY OF THE REACTION BETWEEN FERRIMYOGLOBIN AND HYDROGEN PEROXIDE

Philip George and D. H. Irvine*

Department of Colloid Science, The University, Cambridge, England

Received March 1, 1956

~kBSTRACT

The reaction between ferrimyoglobin, Fe+b(tt20), and hydrogen peroxide results in a one-equivalent oxidation of the hemoprotein giving a higher oxidation state of the molecule whose structure is best represented as a complex ferryl ion, FeMsO. A kinetic study of the formation of this compound, ferrylmyoglobin, shows that there are two reaction paths, the acidic form of ferrimyoglobin, Fe+b(H~O), reacting with H~O2, and either the alkaline form, Fe~bOH, reacting with H202, or the acidic form reacting with O2H-.

Two possible reaction mechanisms for the formation of ferrylmyoglobin are ex- amined in the light of the quantitative kinetic data obtained. The first involves the production of an OH radical in a reaction similar to that usually written for the fer- rous ion reaction, except that the oxidation states are all higher by 1 in the ferrimyo- globin reaction. The second involves the initial formation of a +5 oxidation state in the rate-determining step, which then undergoes extremely rapid reduction to ferryl- myoglobin by reducing groups present in the system (i.e., in the protein part of the molecule). In the lat ter mechanism, ferrimyoglobin would conform to the same pattern of oxidation-reduction behavior shown by the hemoprotein enzymes, per- oxidase and eatalase.

INTRODUCTION

F e r r i m y o g l o b i n is a h e m o p r o t e i n in which the p ros the t i c group, ferr ic p r o t o p o r p h y r i n , is jo ined to a p ro t e in of molecu la r weight 17,000. W h e n H20~ is a d d e d to i t a red c o m p o u n d is fo rmed hav ing wel l -def ined spec t ro - scopic charac te r i s t i c s (1). Th i s compound was or ig ina l ly t h o u g h t to be a complex of the e n z y m e - s u b s t r a t e type , and the v iew was s t r eng thened b y the work of Chance (2), who showed t h a t the k ine t ics of the reac t ions of H~O~ and a lky l hyd rope rox ides wi th perox idase and ca ta lase , hemo- p ro te ins s imi lar in s t ruc tu re to fe r r imyoglobin , could be sa t i s f ac to r i ly exp la ined us ing a modi f i ca t ion of the M i c h a e l i s - M e n t e n m e c h a n i s m of enzyme act ion.

I n prev ious pape r s (3-6) we showed conclus ive ly t h a t the red compound is n o t an e n z y m e - s u b s t r a t e complex b u t a h igher ox ida t ion s t a t e of ferr i -

*Present addresses: P.G. Department of Chemistry, The University of Pennsyl- vania, Philadelphia. D.H.I. Department of Chemistry, University College, Ibadan, Nigeria.

327

3 2 8 PHILIP GEORGE AND n . I-I. IRVINE

myoglobin, one oxidation equivalent above ferrimyoglobin, capable of being formed by other oxidizing agents besides peroxides. In other words, it reacts as if the iron is in the oxidation state -k4. Other experimental studies showed the site of oxidation to be the iron atom itself, and not some other point in the very complicated molecule. Since it represents the first well-authenticated case of a stable derivative of quadrivalent iron, it is of interest in inorganic chemistry; and the reaction by which it is formed with hydrogen peroxide is of significance from the physicochemical standpoint for the following reason.

Our earlier investigations showed that in the reaction of ferrimyoglobin with H202 a transient oxidizing entity was produced along with the stable red compound (3). This entity has many of the properties of the OH radical,

0

Protein FIG; 1.

and the chemistry of the ferrimyoglobin-H202 reaction can be satisfactorily accounted for on the assumption that it is formed:

Fe~+h (H20) -~ H202 --~ FeIM~ v -~ OH + etc. [1]

where Fe~Jb (H20) denotes ferrimyoglobin and F e ~ the higher oxidation state. On this basis, therefore, there is a formal similarity to the first step in the reaction between the ferrous ion and hydrogen peroxide when it is written in the form

2+ 3+ Feaq-~ HeO2 --+ Feaq-~- O H -]- O H - , [2]

the chief difference being that in the hemoprotein case the oxidation numbers are higher by 1 throughout.

Measurements of the change in pH which accompanies the formation of the higher oxidation state by potassium chloriridate, and also of the pH variation of the equilibrium constant for this reaction, show 2 moles of H + to be liberated per mole of ferrimyoglobin oxidized (6), i.e.,

2- IV Fe~+b (H,O) + IrCl~ ~- FeMb-~ 2H ~- IrCl~- zr etc. [31

The structure which best fits these data, in relation to other reactions of the higher oxidation state, is that of a coordination compound of the

FERRIMYOGLOBIN AND HYDROGEN PEROXIDE 329

hypothetical ferryl ion, FeO 2+, with four of the six bonds about the iron atom within the octahedral complex to pyrrolic nitrogen atoms of the porphyrin ring, the fifth linking the iron to a group in the protein, and the sixth binding the oxygen atom which characterizes the ferryl ion type of structure (Fig. 1). The higher oxidation state is thus conveniently referred to as "ferrylmyoglobin."

The present paper deals with the kinetics of its formation by H202 under various conditions of pH, temperature, and ionic strength. The significance of the results, and in particular the energetics, are discussed in the light of previous knowledge of the chemistry of the reaction and similar reactions of peroxidase and catalase.

EXPERIMENTAL

Materials. Ferrimyoglobin was prepared according to the method de- scribed previously (3). The H~O2 was pure 97 % (w/w) free from inhibitors, kindly supplied by Laporte Chemicals Ltd.

Measurement of Rates. The rate of formation of ferrylmyoglobin was determined by measuring the change of optical density with time at a suitable wavelength (409 m~), using a Unicam Quartz spectrophotometer adapted for work at constant temperature.

RESULTS

1. The Effect of pH on the Rate of Formation

Experiments were restricted to the pH region between 8.0 and 9.5, where it was known from previous investigations that only slight destruction of the hemoprotein molecule occurred during the oxidation (3). Within this region good second-order plots were obtained up to 80 % reaction, but beyond this deviations from linearity were noticed. This was to be expected since, through side reactions, a little more than 1 mole of H202 is required to give 1 mole of ferrylmyoglobin. Figure 2 shows some of the ~econd-order plots obtained at 22.0°C. and I = 0.02, and in Table I the values of the second-order constant, kob~., are listed for different pH values at four temperatures.

Inspection of the data shows that /~0b~. decreases slightly as the pH increases from about 8.0 to 8.5, but more markedly in the range 8.5 to 9.5. This effect could arise through the ionization of the peroxide molecule or a group in ferrimyoglobin itself, which is known to have a pK of about 9.0 (7). This group is believed to be a water molecule coordinated to the iron atom in the acidic form, for the ionization is accompanied by a profound alteration in the spectra of the ferrimyoglobin comparable to that resulting from the formation of its cyanide and azide complexes. In these reactions the change in magnetic susceptibility, from that appropriate to a "pre- dominantly ionic" structure to that of a "predominantly covalent" struc-

0 9

0-8

0-5

I I 0-7

o

o 0.6 ._J

I I 5 I0

T IME (rain.)

~ 3 0 PHILIP GEORGE AND D. H. IRVINE

i I 0 15 2 0

FIG. 2. Second-o rde r p lo t s for t h e f e r r i m y o g l o b i n - H 2 0 : r e a c t i o n a t 22°C. a n d I = 0.02 a t p H 8.02 : • ; 8.43 : Q ; 8.81 : [] ; a n d 9.26: ~ .

F e r r i m y o g l o b i n c o n c e n t r a t i o n , a = 4.34 X 10-6M; H:O2 c o n c e n t r a t i o n , b = 11.49 X 10-6M; F e r r y l m y o g l o b i n c o n c e n t r a t i o n , x a t t i m e t.

T A B L E I

Var ia t i on w i th p H of the Observed Second-Order Ve loc i ty Constants (kobs., l. mole -~ sec. -1) for the React ion of H202 wi th Ferr imyoglob in at I = 0.02

(a) T = 1~.0 4- 0.1 (b) T = 22.0 -4- 0.2

pH hobs. pH hobs. 8.07 121 =t= 2 8 .02 203 ± 5 8.27 106 ~ 2 8 .22 183 -4- 6 8 .48 100 ::h 3 8 .43 170 ~ 3 8.66 92 -4- 2 8.61 165 =i= 3 8.86 76 =t= 2 8.81 143 =t= 2 9 .08 68 :i= 2 9 .03 118 -4- 4

9 .26 98 ± 2 9 .46 85 ± 3

(c) T = 29.7 ~ 0.3 (d) T -- 36.0 -4- 0.4 pH hobs. pH hobs.

8 .38 282 =E 11 8 .34 473 ::t:: 11 8.56 251 ~ 6 8 .52 413 -4- 11 8.76 224 =t= 7 8.72 335 =t= 9 8.98 211 ~ 7 8 .94 311 4- 12 9.21 181 ~= 11 9.17 277 4- 9

FERRIMYOGLOBIN AND HYDROGEN PEROXIDE 3~1

ture, leaves little doubt that the cyanide and azide are bonded to the iron. The ionization in question can thus be represented by the equation

Fe +(H20 ) KF~ Fe~bOH + H + [4]

with the acidic form carrying a net charge of + 1, since two of the three positive charges of the ferric iron are cancelled by negative charges on two of the pyrrolic nitrogen atoms of the porphyrin ring. The correspond- ing formula for ferrylmyoglobin is FeMbO.

If the peroxide reacted exclusively through its anion, O2H-, then in the pH range 8.0 to 8.5, kob~. would be approximately proportional to 1/H +. The data in Table 1 are clearly inconsistent with this mechanism. But the two following mechanisms give identical kinetic expressions which account very well for the observed variation. First, if acidic and alkaline ferrimyo- globin react with peroxide with velocity constants/ca and kb

Fe+b (H~O) + H20.~ ~ FeMbO + etc.

FeMbOH + H202 kb * Fe~bO + etc.,

]Cob~. would vary with hydrogen ion concentration according to the equa- tion

/~,H + kbK~e ]Cobs" KFe + H + + KFe + ~ H +' [5]

where KF~ is the ionization constant. This may be rearranged to give

kobS.(KF~ + H +) = /~H + + kbKF~. [5a]

On the other hand, if acidic ferrimyoglobin exclusively reacts with H~O2 and O2H-, with the velocity constants k~ and k'b,

Fe+b (H~O) + H202 k~ , FeMbO + etc.

Fe+b (H20) + O2H- k~ ~ FeMbO + etc.

then the rate equation is given by

H+ F kaH+ ' k'bK, ~ k°bs" = (gFe + H+) " LKp+H+ ~- Kp + H+J [6]

where K~ is the ionization constant of hydrogen peroxide. Since K~ 2 X 10 -12, for the range of hydrogen ion concentration employed, H + >> K~, so this equation can be simplified and rearranged to give

kobs.(KFe + H +) = k~H + + /c~Kp, [6a]

which is identical in form with Eq. [5a]. With the use of the values of KF, obtained from the work of George and Hanania (7), the resulting plot of

2 0

0

%

X 16

x +

~. 12

D

~¢ 8

I I I I I 2 4 6 8 I 0

332 PHILIP GEORGE AND D. H, IRVINE

H + X 10 9

FIo. 3. Plot of kob,. (K~o + H +) against H + for data obtained at 22°C. and I = 0.02. KFe is the ionization constant, Fe+b (H20) ~ Fe+bOH + tt +.

TABLE I I Variation with Temperature of the Velocity Constants ka and kb(l. mole -1 sec. -1) for

the Reaction of H202 with the Acidic and Alkaline Forms of Ferrimyoglobin T, °C ka kb

15.0 ± 0.1 122 ± 2 22.0 ± 1.5 22.0 -4- 0.2 220 ± 6 49.0 ± 3.0 29.7 ± 0.3 320 ± 16 140 ± 15 36.0 ± 0.4 607 ± 33 197 ± 22

koUs. (K~e + H +) against H + is linear, as shown for the data at 22.0°C. and I = 0.02 in Fig. 3. Hence on this analysis alone both mechanisms are equally valid. Values of ka have been obtained f rom the slopes of the lines at the four temperatures used, but values of kb were arrived a t by plott ing kobs. (KFo ~ H + ) / H + against 1 / H +. This method was Used in preference to tha t of utilizing the intercepts on the previous plots because, al though subject to large errors, it was much more sensitive. The values of ka and kb are recorded in Table I I , and f rom the lat ter k'b can be evaluated f rom the relationship kbKFe = k'~K~ which follows f rom the comparison of Eqs. [5a] and [6a].

I t m a y be noted tha t in the plot of kobs. (KFo -~ H +) against H + in Fig. 3, the point corresponding to the most acidic solution p H 8.02, does not lie on the straight line, and tha t the discrepancy is outside the limit of

F E R R I M Y O G L O B I N A N D H Y D R O G E N P E R O X I D E 333

experimental error. Now it is known that there is another ionizable group in ferrimyoglobin, possibly on the amino acid residue to which the ferric protoporphyrin is attached, whose ionization affects the velocity constants, and hence the equilibrium constants for the formation of ferrimyoglobin fluoride, and cyanide (8-10). It has a pK of about 6.1 at 20°C. This ioniza- tion, therefore, is very likely the cause of the discrepancy. Dalziel and 0'Brien (11) have observed a similar effect in the reaction of H202 with ferrihemoglobin, a hemoprotein of molecular weight 68,000 with four ferric protoporphyrin groups attached to it, which also possesses similar "heine- linked" ionizable groups (12). Unfortunately, in solutions acid to pH 8.0, since side reactions occur involving irreversible oxidation and destruction of the protoporphyrin (3), a reliable quantitative investigation of the participation of these ionizing groups cannot be made.

2. The Effect of Temperature

From plots of log10 /c~ and loglo kb against 1/T the activation energies were evaluated, and together with data for the variation of KFo and K~ with temperature (7, 13), the activation energy for the reaction with velocity constant krb was calculated, giving

/ ~ = 10 n '~±° '7 e x p . - - ( 1 2 , 6 0 0 ~ IO00)/RT kb = 1016"5~2"2 e x p . - - ( 2 0 , 0 0 0 ~ : 3000)/RT k~b = 10 I7"4±2"2 e x p . - - ( 1 7 , 6 0 0 : h 2500)/RT

3. The Effect of Ionic Strength

A direct experimental determination of the effect of ionic strength on the rate constant kob,. was impracticable because changing the ionic strength changed the pH of the buffer solution. It was therefore decided to estimate the effect by measuring the rate constant at a known ionic strength, determining the pH of the solution, and making a comparison with the constant calculated from Eqs. [5a] and [6a] according to various ionic strength variations of k~ and kb, of k~ and/~'b. With the spectrophoto- metric technique employed it was not possible to carry out experiments in the absence of air, and absorption of carbon dioxide by very dilute buffer solutions in the pH range about 9.0 prevented an investigation of the ionic strength effect in this particular range, where the term containing kb or k'b makes a significant contribution to kob,.. Data obtained at a pH of about 8.2 are given in Table III. On the assumption that k,, and kb or k'b, are independent of ionic strength, calculated values of kob~. agree quite well with the experimental values, as shown by the comparison of columns 3 and 4. Since at this pH the term containing/~b or k'b contributes only about 8 % of/Cob,., the variations of kob,. with ionic strength would in any case be very little affected by any variation of kb or k'b, and it may be concluded that k~, which contributes the remaining 92 %, is therefore inde-

3 ~ P H I L I P G E O R G E AND D. H. I R V I N E

TABLE III The Effect of Ionic Strength on the Rate Constant at T = 25.0°C.

I pI-I kobs. k~al~.

0.16 8.10 240 :i: 6 235 ::h 8 0.04 8.18 230 ± 5 227 :i: 8 0.02 8.21 225 ± 5 224 :h 8 0.01 8.24 230 ± 10 222 ± 8 0.005 8.25 250 ::h 5 222 ::h 8 0.0025 8.23 245 ::h 5 226 ~ 8

pendent of ionic strength within experimental error. This is good evidence that ka does not correspond to a reaction between ionic species, and it therefore supports the conclusion drawn from the kinetic analysis tha t the reaction involves the undissociated H202 molecule.

4. The Effect of Added Reducing Agents on the Rate of the Reaction

Previous experiments had shown that if a reducing agent was added to ferrimyogtobin before the addition of H2Q, then, provided the concentra- tion of reducing agent did not exceed that of the ferrimyoglobin, ferryl- myoglobin was formed to almost the same extent as in the absence of the reducing agent (3). Measurement of the rate of formation under these conditions showed that this was also unaffected by the presence of the following reducing agents, ferrocyanide and iodide ions, tartaric acid, and mannitol. In these reactions deviations from second-order kinetics were observed after about 60 % reaction. This can be at t r ibuted to partial reduction of the higher oxidation state by the reducing agent. When excess reducing agent was used, however, the couple ferrylmyoglobin/ferrimyo- globin acted as a catalyst for the oxidation of the reducing agent by per- oxides. An account of this will be published elsewhere.

Discussion

The difficulty in making a choice between two reaction mechanisms that lead to the same dependence of velocity constant on hydrogen ion concentration is not unfamiliar in reaction kinetics. In the ionic-iron- hydrogen peroxide system, the electron transfer Fe ~+ + 02- -+ Fe 2+ + 02 is preferred to the hydrogen atom transfer FeOH 2+ + HO~ --~ Fe z+ + H20 + 02 from a comparison with corresponding reactions of the cupric ion (14). In the oxidation of hydroqninones by ferric ions the electron transfer step Fe ~+ + Q- --~ Fe z + + Q is preferred to hydrogen atom trans- fer in FeOH ~+ + QH --> Fe 2+ + H20 + Q in view of the magnitudes of the entropy of activation and the overall entropy change (15). In the present case no choice can yet be made between FeMbOH reacting with H202, and Fe+b (H~O) reacting with O2H-; moreover in the case of such hemoprotein reactions a special difficulty can arise as the following con- siderations show.

FERRIMYOGLOBIN AND HYDROGEN PEROXIDE ~ 3 5

The values of the temperature-independent factors for these two reac- tions are 10 ~6.5~-~ and 1017.4±2-2, respectively. High values are more ap- propriate to reactions between oppositely charged ions (16), so this cri- terion appears to favor the latter reaction, viz., Fe+b (H20) + O2H-. However, in the pH range 8 to 9.5 where the measurements were made, which is alkaline with respect to the isoelectric point of myoglobin, studies of the variation of the ionization constant, KFo, as a function of ionic strength, indicate that the ionization is accompanied by a change in charge on the iron atom of - 2 to - 3 , provided the simple Debye-Htickel theory is applicable to these systems (7). In other words, the net negative charge on the protein as a whole, which titration studies show is of the order - 1 0 in this pH range (6), results in the iron atom "carrying" an effective negative charge of - 2 in its acidic form and - 3 in its alkaline form. This being so, the reaction of Fe+b (H20) with O2H- would involve species with charges of - 2 and - 1 and so the magnitude of the tempera- ture:independent factor really favors the other alternative, FeMbOH react- ing with H20~.

There is little doubt, however, that the chief reaction path in the pH range 8 to 9 entails Fe+b (H~O) reacting with H202, and the heat of the reaction can be calculated using thermochemical and bond energy data from the literature to compare with the observed activation energy of 12.6 4- 1 kcal./mole. If it is assumed, as the evidence would appear to indicate, that an OH radical is formed together with ferrylmyoglobin in the oxidation by H20~, and that one H + is liberated, the first step in the reaction of H202 with acidic ferrimyoglobin would be given by the equation:

FeM+b (It~O) + H202 --* FeMbO + H30 + + OH. [7]

The heat of this reaction can be calculated from the heat of the following steps (6):

FeM+b (H20) + IrCl~ --* FeMbO + 2H + + IrCl~- - 9.0 4- 1.0 kcal./mole [8]

IrCl~- + H + --~ IrCl~ + ~H2 - 30.5 4- 0.5 kcal./mole [9]

H202 + 1/~ H~ ~ OH. + H:O + 21.6 4- 2.0 kcal./mole [10]

to give the value --17.9 5= 3.5 kcal./mole. The heat of reaction [10] is 22.3 kcal./mole according to the standard heats of formation quoted by Latimer (17) for (H~O)I, (H20~)~q. and (OH.)~ together with the values of 10.5 and 1.4 kcal./mole for the heats of solution of OH and H2, respectively, adopted by Evans, Hush, and Uri (18). From the bond energy data com- piled by these authors, the heat of reaction [10] is found to be 20.8 keal./ mole. A mean value of 21.6 has been taken as the best approximation and the uncertainty put at ~2.0 kcal./mole.

The endothermicity of reaction [7] is thus 17.9 5= 3.5 kcal./mole, and the observed activation energy 12.6 4- 1.0 kcal./mole. Although the

~36 PHILIP GEORGE AND D. n. IRVINE

discrepancy is quite small, amounting to only 0.8 kcal./mole if the highest activation energy and lowest endothermicity are chosen, these calculations suggest that reaction [7] may not be the actual rate-determining step. Con- sideration should therefore be given to a second reaction mechanism which could satisfy the energetics and still be consistent with the other experi- mental observations. This entails a rate-determining step in which peroxide effects a two-equivalent oxidation, giving a compound in which the iron is in the oxidation state +5, followed by the very rapid reduction of this compound to ferrylmyoglobin by reducing matter present in the system.

rate-determining FeVb + etc. Fe+ b (H20) + H~O~ step

FeVb single-equiv, reduction ~ Fe~bO + etc.

The assumption of a rapid reduction step is essential to account for the fact that the rate of formation of ferrylmyoglobin is unaffected by the addition of reducing agents, and for the observation that during the course of the reaction only the two species, ferrimyoglobin and ferrylmyoglobin, have been detected spectrophotometrically. On this hypothesis, the reac- tions previously attributed to the OH radical would be assigned to the single-equivalent oxidation products formed by the attack of the + 5 oxida- tion state on the reducing matter present, e.g., groups on the protein. Transient radical species could easily result from an attack of this kind. The reaction which was taken to indicate the presence of the OH radical, namely, the luminol test, would also occur with these other species, pro- vided they were strong enough oxidizing agents. The complex ion, Ru (dipy)3 ~+, with E0 = 1.35 v., and the permanganate ion, E0 = 0.56 v., were found to oxidize luminol with an intense luminescence, showing this test not to be a characteristic of the OH radical as such, but merely an indication of its great oxidizing power and the rapidity with which it reacts.

In this connection it is interesting to note that Cahill and Taube (19) have recently presented evidence which suggests that in the ferrous ion- hydrogen peroxide system, two-equivalent oxidation to give a quadrivalent iron compound is the important initiating reaction

Fe~ + + H202--* Felq v. + etc.

However, the detailed kinetics have been explained so far only by a free radical mechanism in which the OH. radical is formed initially as in re- action [2]: the need to postulate the formation of a quadrivalent iron com- pound, possibly the ferryl ion, arose only when high ferric ion concentra- tions were present (14, 20).

An interesting feature of the ferrimyoglobin-hydrogen peroxide reaction is the contrast between the rate data for the two reaction paths. In the reaction of H202 with the acidic form, the activation energy 12.6 ~ 1.0 kcal./mole is relatively low, favoring the reaction, whereas the entropy

F E R R I M Y O G L O B I N A N D H Y D R O G E N P E R O X I D E

'%. H - ~

337

~.__~ i F , ~ O - - ~ b' , , ~, Mb ' ~ , ~ ~

FIG. 4.

of activation, calculated to be -6 .6 ± 1.5 e.u., is unfavorable. In the case of the alkaline form reacting with H202, or alternatively the acidic form reacting with O:H-, the activation energies, 20.0 ± 3.0 and 17.6 ± 2.5 kcal./mole are much higher, but the entropies of activation are favorable, +15 ± 10 and + 19.5 ± 10 e.u., respectively. Since the formation of ferrylmyoglobin requires the removal of both hydrogen atoms bonded to the oxygen in the coordinated water molecule of acidic ferrimyoglobin, or the single hydrogen atom attached in the OH group of the alkaline form, the various transition states can be pictured as shown in Fig. 4, for the mechanism involving OH radical production. According to the second mechanism involving the initial formation of a +5 oxidation state, pro- vided this also had a bonded oxygen atom structure, the required align- ment of the reactants would be substantially the same. The negative entropy of activation of reaction a could be attributed to the special orientation required for the H202 molecule, with the peroxidic oxygen atoms bridging the hydrogen atoms of the coordinated water molecule. In view of the favorable entropy of activation for the second path, the need for no special orientation in reaction b, in contrast to reaction b' which resembles reaction a in this respect, could be regarded as evidence that this is the more likely of the two alternatives. The differences in acti- vation energy agree with this interpretation. I t is likely that in ferrylmyo- globin a strong Fe--O bond is formed, and this could also be the case in the +5 state. But there is already quite a strong Fc O bond in the alkaline form of ferrimyoglobin, since kinetic studies show this form to be unreactive in cyanide and fluoride complex formation (9), whereas, when the acidic form reacts, the formation of the Fe--O bond in the higher oxidation state would provide a greater balance of energy to offset the loss from bond- fission.

Finally a comparison may be drawn between the reaction of peroxides with ferrimyoglobin and with the enzymes peroxidase and catalase, two

338 PHILIP GEORGE AND D. H. IRVINE

similar hemoproteins in so far as they have the same prosthetic group. In both of the enzyme systems the participation of two intermediate com- pounds is well established. One of these resembles ferrylmyoglobin, in that it is a single-equivalent oxidation product; the other is either a two- equivalent oxidation product, or a peroxide complex, i.e., containing H202 or O~H- as the ligand (21). Recent studies (22) favor the + 5 oxidation state, and the interrelationship of the various compounds can be expressed by the following scheme, where Fe~ + denotes the ferriprotoporphyrin pros- thetic group of peroxidase or catalase, and AH~ denotes a reducing agent,

2-equiv. Fe 3+ --~ H202 oxidation ~ Fe v -F etc.

one-equiv. [Fe~V.AH ] separation FeipV + A H + etc . Fe v + AH2 reduction )

$ $+Am

Fe3p + + A + etc. Fe3p + + etc. Catalase Action Peroxidase Action

The catalytic action of peroxidase involves, in turn, the two single-equiva- • V ~-~ IV i ~ IV ~+ lent reductlon steps, Fep --~ ~ep ann ~e, -+ Fe, (23). Although theaction

of catalase appears to involve a distinct two-equivalent reduction, Fe v --~ 3+ Fep , it is more likely, for two reasons, that a single-equivalent reduction

of Fe v occurs initially and then a further reduction of Fe~ v within the reaction "cage," a small fraction escaping during each cycle. Firstly, the Fe~ v compound of catalase is actually formed progressively during the course of the oxidation of those reducing agents which appear to undergo the two-equivalent oxidation (24), and secondly, in the presence of the appropriate reducing agent (aromatic amines or phenols), catalase action, just ]ike that of peroxidase (25), proceeds by the two distinct single- equivalent reduction steps• If the ferrimyoglobin reaction gives a + 5 oxida- tion state, the reaction of all three hemoproteins would fit into the same pattern of oxidation-reduction behavior, the differences being of degree rather than kind. The greater catalytic activity of peroxidase and catalase would then be attributable primarily to the higher speed of the initial two-equivalent oxidation of the prosthetic group. The indiscriminate re- duction of the + 5 state of myoglobin to ferry]myoglobin would also dis- tinguish the two systems, through the waste of oxidizing capacity. On the other hand, if the ferrimyoglobin reaction proceeds with OH radical forma- tion, then the smaller catalytic activity would be attributable essentially to the lower speed of the single-equivalent oxidation of the ferrimyoglobin.

I:~EFERERrCES

1. KEILIN, D . , AND HARTREE, E . F . , Proc. Roy. Soc. (London) Bl17, 1 (1935). 2. CHANCE, B. , " T h e E n z y m e s , " Vol. 2, P a r t 1, p. 428. Academic Press , N e w York ,

1951.

FERRIMYOGLOBIN AND HYDROGEN PEROXIDE 339

3. GEORGE, P., AND IRVINE, D. H., Biochem. J. 52,511 (1952). 4. GEORGE, P., AND IRVINE, D. H., Biochem. J. 55,230 (1953). 5. GEORGE, P., AND IRVlNE, D. I-I., Biochem. J. 58, 188 (1954). 6. GEORGE, P., AND IRVINE, D. I'i., Biochem. J . 60,596 (1955). 7. GEORGE, P., AND HANANIA, G. I. ]=[., Biochem. J. 52,517 (1952). 8. GEORGE, P., AND HANANIA, G. I. i . , Biochem. J. 56, xxxvii (1954). 9. GEORGE, P., AND HANANIA, G. I. i . , Biochem. J. 56, xxxviii (1954).

10. GEORGE, P., AND HANANIA, G. I. H., Discussions Faraday Soc. No. 20, 216 (1955). 11. DALZIEL, I~., AND O'BRIEN, J. R. P., Biochem. J. 56, 648 (1954). 12. WYMAN, J., JR., Advances in Protein Chem. 4,410 (1948). 13. EVANS, M. G., AND URI, N., Trans. Faraday Soc. 45,224 (1949). 14. BARB, W., BAXENDALE, J. H., GEORGE, P., AND HARGRAVE, I~. R., Trans. Faraday

Soc. 47,462 (1951). 15. BAXENDALE, J. I~., I'IARDY, I'i. l:~., AND SUTCLIFFE, L. H., Trans. Faraday Soc. 47,

963 (1951). 16. GLASSTONE, S., LAIDLER, K. J., AND EYRING, H.,"The Theory of Rate Processes."

McGraw-Hill, New York, 1941. 17. LATIMER, W. M., "Oxidation Potentials." Prentice-Hall, New York, 1952. 18. EVANS, M. G., HUSH, I~. S., AND URI, N., Quart. Revs. (London) 6, 186 (1952). 19. CAHILL, A. E., AND TAUBE, H., J. Am. Chem. Soc. 74, 2312 (1952). 20. BARE, W., BAXENDALE, J. H., GEORGE, P., AND HARGRAVE, K. R., Trans. Fara-

day Soc. 47,591 (1951). 21. GEORGE, P., AND IRVINE, D. ~Lt., Brit. J. Radiol. 27,131 (1954). 22. GEORGE, P., Science 117, 220 (1953). 23. GEORGE, P., Biochem. J. 54,267 (1953). 24. CHANCE, B., J. Biol. Chem. 179, 1341 (1949). 25. GEORGE, P., Biochem. J. 52, xix (1952).