The Specific Reactions of Iron in Some Hemoproteins

The Specific Reactions of Iron in Some Hemoproteins PHILIP GEORGE Department of Colloid Science, University of Cambridge, England

CONTENTS Page .............................................. 367 bin, Myoglobin, Peroxidase, and Catalase.. 369 1. Structure. . ...................................... 369 ematin . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 372 3. Compounds of Hemoglobin, Peroxidase, and Catal . . . . . . . . . . . 374 111. Hemoglobin and Myoglobin Autoxidation and Other 5 . . . . . . . . . . . 381 IV. Peroxidase Reactions. . . . . . ... V. Catalase Reactions.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 393 1. Catalytic Decomposition of Hydrogen Peroxide. . . . . . . . . . . . . . . . . . . . . . 393 and Alkyl Hydro2. Complexes of Catal . . . . . . . . . . . . . . . . . . 398 peroxides. . . . . . . . . . 3. Catalase in Coupled . . . . . . . . . . . . . . . . . . 400 4. Reaction Mechanisms with the Catalase-Peroxide Complexes. . . . . . . . . . 402 VI. The Mechanism of These Hemoprotein Reactions. . . . . . . . . . . . . . . . . . . . . . 404 1. Free Radical Mechanisms with Ferrous and Ferric Ions.. . 2. An Examination of the Possibility of Similar Free Radical Reactions with the Hemoproteins . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 411 3. The Applicability of These Free Radical Mechanisms in the Hemoprotein Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 415 a. Reactions of Peroxidase and Catalase. . . . . . . . . . . . . . . . . 415 b. Oxidative Reactions of Hemoglobin and Myoglobin Involving Oxygen 420 VII. Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 424 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 425

I. Introduction. ....

I. INTRODUCTION I n the editorial preface to the first volume of Advances in Catalysis the decision was made known not to publish reviews of specialized topics in biocatalysis but from time to time to bring reports in which the relationship and parallelism between this special field and “normal ’’ catalysis are discussed. This is the first of these reports. Its purpose is to examine the reactions of four hemoproteins, hemoglobin, myoglobin, peroxidase, and catalase, which all contain the same coordination compound of iron-ferrous or ferric protoporphyrin attached to different protein molecules, with oxygen, hydrogen peroxide, and in a few cases additional reducing substances. Some of these reactions are specific; 267

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PHILIP GEORGE

that is, they are either more highly developed, or developed to an extent that they can be regarded as unique reactions, when compared with the corresponding reactions of the other three hemoproteins or the reactions of ionic iron. These specific reactions may be listed as follows: Hemoglobin and Myoglobin. Ferrous compounds which combine reversibly with molecular oxygen, e.g., Hb

+ 02

HbOz

Both association and dissociation reactions are extremely rapid. The oxidation to the corresponding ferric compound is in comparison a very slow reaction. Perozidase. A ferric compound which utilizes hydrogen peroxide to oxidize reducing substances according to the general ;scheme:

+ +

Peroxidase HzOl -+ Peroxidase . [H&] Peroxidase [H202] AH1 -+ Peroxidase A -t2Hz0

+

which has been named “peroxidatic activity.” Peroxidase can be reduced by strong reducing agents to the ferrous compound. Catakse. A ferric compound whose biological function may be the oxidation of alcohols or other substrates by hydrogen peroxide in an analogous reaction to peroxidase and which is the most efficient catalyst known for the decomposition of hydrogen peroxide into water and molecular oxygen at room temperature. This has often been called “catalatic” activity. It is a most remarkable feature of catalase that its ferrous form cannot be prepared directly using strong reducing agents or electrolytic reduction. A ferrous compound is formed in the presence of hydrogen peroxide and azide ions, but its relation to the true ferrous form or whether in fact it is the ferrous form corresponding to the ferric form of the free enzyme is not known. Hemoglobin and myoglobin in their ferric forms show rudimentary peroxidatic and catalatic activity, but ferrous peroxidase does not combine reversibly with molecular oxygen. Ionic iron also gives the hydrogen peroxide reactions but not the combination with oxygen. Many experimental studies now support a free radical mechanism for these reactions of ionic iron, and it is the main theme of the report to examine the data on the hemoprotein reactions to see whether similar reaction mechanisms are first of all possible and if so whether there is evidence that the reactions do proceed in $his way. It is unavoidable that many topics of great biochemical importance and physico-chemical interest must be omitted in a, report of this kind. Two recent reviews by Wyman (1) and Theorell (2), articles in Respiratory Enzymes (3) and Haemoglobin (4) by various authors, and the

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

369

recent monograph Haematin Compounds and Bile Pigments by Lemberg and Legge ( 5 ) deal in great detail with these other themes. I n addition these sources review the very extensive biochemical investigations on the hemoprotein reactions discussed in this report and so reference is often made to these sources where parallel references may be found. A short section summarizing the relevant properties and reactions of iron protoporphyrins and the four hemoproteins precedes the detailed account of their reactions with hydrogen peroxide and molecular oxygen.

11. GENERALCHEMISTRY OF HEMOGLOBIN, MYOGLOBIN, PEROXIDASE, AND CATALASE 1. Structure

All four hemoproteins have the same reactive group, iron protoporphyrin :

where R = -CH8

;J-fR R

qN.

R Ri z

7 \

R2

= -CH&HzCOOH -CH=CHz

R2

The assignment of the double and single bonds in this figure is purely formal, and it is well established from X-ray studies of the structurally similar phthalocyanines that these molecules are resonance hybrids. Even in free phthalocyanine where two of the four pyrrole nitrogen atoms carry hydrogen atoms the interatomic distances in the inner 16-atom ring are almost identical, having values of 1.33 A. or 1.34 A. The porphyrin forms its iron compound by the replacement of the two pyrrole hydrogen atoms. Thus the ferrous derivative, called heme, has no ionic charge. The ferric derivative has a single positive charge; as the chloride it is called hemin and as the hydroxide hematin. It will appear later that the net charge on the iron atom plays an important part in determining what type of addition compound the molecule forms. These formal charges on the iron atom may not necessarily represent the charge on the molecule as a whole, for in the case of the free molecule the ionization of both propionic acid side chains could clearly result in heme carrying two negative charges and the ferric compound being no longer a chloride but a zwitterion carrying an additional negative charge. There is good evidence, however, particularly in the case of peroxidase, that in the hemoproteins these side chains may be concerned in the linkage of the iron protoporphyrin to the protein molecule, and the contribution

370

PHILIP GEORGE

of the iron protoporphyrin to the total charge in the molecule as a whole depends on the precise nature of this linkage. The four hemoproteins, hemoglobin, myoglobin, peroxidase, and catalase, differ only in the protein to which the iron protoporphyrin is attached. Molecular weight determinations give values 68,000, 17,000, 44,000, or 40,000 (Keilin and Hartree, 48) and 225,000 respectively. Analytical data show hemoglobin and catalase to have four heme or hematin groups per molecule whereas myoglobin and peroxidase have only one. The recent experimental data on the size and shape of these hemoproteins has been reviewed by Wyman (1). Not only do the four hemoproteins differ in molecular weight but also in the relative proportions of their constituent amino acids. Only for. hemoglobin and myoglobin are nearly complete analytical data available. Tristram (4) has recently summarized the results of many different workers. Both proteins are made up of about twenty different amino acids. In units of weight per cent, leucine and isoleucine 16.:3%, glutamic acid 16.48%) lysine 15.5%) histidine 8.50/,, aspartic acid 8.2% and alanine 7.95% predominate in myoglobin, six amino acids accounting for 73.4% of the molecule. Hemoglobin has a rather more varied structure; eight amino acids account for 75.75% of the molecule: leucine 15.4%) aspartic acid 10.6%) valine 9.1%, histidine 8.7%, lysine 8.5%) glutamic acid 8.15%, phenyl alanine 7.9%, and alanine 7.4%. In one respect there is a marked difference between these two proteins: myoglobin contains no cystine or cysteine, hemoglobin contains enough t o give a maximum of three disulfide links per molecule, but since -SH groups are known to be present there cannot actually be more than two disulfide links. The analytical data obtained by Theorell and Akeson (6) for peroxidase and catalase respectively enabled them to calculate the lysine, arginine, and histidine contents (see also Bonnichsen, 7). These results are summarized in Table I together with the relevant values for hemoglobin and myoglobin in terms of amino acid per molecule (R/M), and the TABLE I Typical Amino Acid Contents

Hemoprotein

Histidine R/M R/68,000

A rginine

R/M

R/68,000

2 14 18 27

8 14 28 8

R/M ~

Myoglobin Hemoglobin Peroxidase Catalase

9 36 3 15-16

36 36 5 5

Lysine R/68,000

~~

18 38 12 32 ~~

Note. R / M Number of residues per molecule. R/68,000 Number of residues assuming molecular weights all to be 68,000.

72 38 19 10

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

371

equivalent number of residues if the molecular weight of all four hemoproteins was 68,000 to facilitate comparison (R/68,000). The total number of residues for a molecular weight of 68,000 is of the order of 580. These data illustrate very well the differences in composition. Histidine and lysine are present in hemoglobin and myoglobin in relatively high amounts but not in peroxidase and catalase. The arginine content of peroxidase is higher than in the other three proteins. A further distinguishing feature of peroxidase is that it contains a high percentage (18.4% of insoluble hydrolyzate) of substances which are not amino acids (Theorell and Akeson, 6). These have been tentatively identified with acid carbohydrate molecules in accord with the observation that peroxidase gives a positive result in the Molisch reaction, and that the isoelectric point is lower than would be expected from a predominance of basic amino acids. X-ray studies of hemoglobin and myoglobin have yielded very valuable information on the size and shape of these molecules. Reference may be made t o the review articles by Perutz (4) and Kendrew (4). The hemoglobin molecule can be pictured in an idealized form as a cylinder of 57 A. diameter and 34 A. in height containing four layers of scattering matter. Myoglobin is similar in shape and dimensions to one of the four layers of hemoglobin: it consists of a disk about 9 A. thick with other dimensions not greater than 57 A., the most probable values being 51.5 by 37 A. The Link between the Prosthetic Group and the Protein. I n all the four hemoproteins, hemoglobin, myoglobin, peroxidase, and catalase, the iron protoporphyrin group is joined to the protein by a coordinate bond to the iron atom. This may be regarded as occupying the fifth coordination position about the iron atom which then binds a water molecule or OH group to complete a stable octahedral coordination unit (Keilin and Hartree, 112). There is some evidence that the protoporphyrin side chains are also involved for not all iron porphyrins combine with the free protein when obtainable to give active compounds. Examples are given below. The observation that hemoglobin is split into heme and free globin in weakly acid solution and the fact that globin contains a high proportion of basic amino acids suggests that the hemes are bound to basic groups in the protein. This reaction is reversible and the reconstituted hemoglobin shows all the properties of the original hemoglobin except for slight spectroscopic differences (Jope, 4). Other ferrous porphyrins have been found to combine with free globin to give compounds which still give the reversible oxygenation reaction : These are mesoheme, hematoheme, and diacetyldeuteroheme containing ethyl, hydroxy ethyl

372

PHILIP OEORQE

and acetyl groups in place of the vinyl side chains of ferrous protoporphyrin. Further discussion of this topic is beyond the scope of this review for it involves the effect of pH and carbon dioxide on the hemoglobin oxygen reaction and the determinations of the ionizable groups on hemoglobin, oxyhemoglobin, and methemoglobin by various titration procedures (Wyman, 1). The results of these studies mainly favor the view that the heme is joined to a histidine group, but a linkage of this type cannot account for the effect of carbon dioxide (Roughton, 8). There is very little evidence on the nature of the linkage in myoglobin and catalase. All attempts to split the hematin group off catalase reversibly have failed. In the case of peroxidase the hematin can be removed by treatment with an acetone-hydrochloric acid mixture a t -5" t o -10°C. and the enzyme reformed by treating the colorless protein with alkaline hematin. Replacing the hematin by deuterohematin or mesohematin which have hydrogen atoms or ethyl groups instead of vinyl groups gave an enzyme with 63% and 57% of the original activity. Hematohematin with a-hydroxy ethyl groups gave no activity (Theorell, 2; see also Theorell and Maehly, 113, and Maehly, 114). Trials with other ferric porphyrins showed that the two propionyl side chains are essential and that these must occupy the positions they do in protoporphyrins. On the basis of differential titration data obtained for peroxidase itnd its free protein Theorell has discussed the possibility that the hematin group is joined to a carboxyl group of the protein (2,6). When heme or hematin are joined t o the various proteins in hemoglobin, peroxidase, and catalase, they combine with a larger range of compounds to give stable octahedral coordination complexes than they do by themselves. To illustrate this the chief derivittives of heme and hematin are listed below before proceeding to describe the hemprotein complexes. 2. Compounds of Heme and Hematin

Both heme and hematin show a marked tendency to aggregate in aqueous solution which has complicated many of the experimental investigations of their reactions. Four of the six valencies of the central iron atom are held in the planar porphyrin ring, and when heme and hematin are in true solution the fifth and sixth bonds completing an octahedral complex are assumed to be occupied by water molecules. These compounds, i.e., heme and hemin, can be represented as H20.Fe,.H20 and H20.Fe,+.H20, where Fe, and Fe,+ represent ferrous and ferric protoporphyrin respectively. Ionization of one of the water molecules attached to heme in solution

T H E SPECIFIC REACTIONS O F I R O N I N SOME HEMOPROTEINS

373

has not been observed, but in the case of hematin a pK of 7.6 has been measured for the ionization (Shack and Clark, 9) : H20,Fe,+Hz0

H20.Fe,.0H

Hemin

+ H+

Hematin

In strongly alkaline solution, however, there is good evidence that heme forms a compound with two hydroxyl groups [H0.Fep.0H]2-, (Keilin, 10). Heme but not hemin or hematin reacts with carbon monoxide giving a compound presumably H20.Fe;CO; it also combines with methyl isocyanide. Hematin but not! heme gives a complex with hydrogen peroxide which Haurowitz represents as H20.xFe;OH (11; see also Haurowitz, Brdicka, and Kraus, 12). A great variety of nitrogenous bases combine with heme including pyridine, nicotine, a-picoline, imidazole and its derivatives e.g. 4 methylimidazole, piperidine, methylamine, and ammonia. These compounds, which can be formulated as B.Fe,.B, are called hemochromogens. Many denatured proteins form hemochromogens particularly denatured globin. One of the molecules of the base can be replaced by carbon monoxide giving the compound B.Fe,CO. Cyanide ions react with heme forming the complex [CN.FeP.CNl2-,and hemin yields [CN.Fe,.CN]-. The same nitrogenous bases that react with heme also react similarly with hemin but have in general a much lower affinity. These compounds are called parahematins and with the corresponding hemochromogens form a well-known redox system. B.Fe,.B

[B.Fe,.B]+

Hemochromogen

+ electron

Parahematin

The potentials of these systems will be discussed later in connection with estimates of the ionization potential of the hemochromogen in aqueous solution. The effect of increasing alkalinity on these systems is somewhat complex for there is evidence in some cases that an equilibrium of the following type can occur. [B.Fe,.B]+

+ OH-

B.Fe,.OH

+B

These topics are discussed in greater detail by Lemberg and Legge (5) and Mansfield Clark (13). Like hemin the parahematins react with hydrogen peroxide. For pyridine parahematin Haurowitz, Brdicka, and Kraus (12) have suggested that the following replacement reactions occur : [Py.Fe,.Py]+Cl-

HzOz

HlOZ

[H202.FeP.Py]+C1--+ [H202~Fe,~Ha02]+C1-

374

PHILIP GEORGE

This behavior may be contrasted with t h a t of the hemochromogens where there is no evidence for similar replacement reactions, but considerable oxidative attack on the porphyrin ring occurs giving compounds named by Lemberg, Cortis-Jones, and Norrie oxyporphyrin hemochrome and verdohemochrome (14).

>

I n the former a methene group ‘CH

/

>COH and in the latter the methene group is replaced by

becomes

0 and an

adjacent carbon atom on a pyrrole ring carries an OH group.

3. Compounds of Hemoglobin, Peroxidase, and Catalase

One of the most characteristic features of these hemoproteins when compared with many other coordination complexes of iron is the very great stability of the protein-iron protoporphyrin unit. I n none of the normal reactions of these hemoproteins, where the protein remains in its native state, is there any evidence of t h e link being broken. The origin of this stability is unknown but would appear t o arise from t h e dissociation process being extremely slow. Indirect evidence of this is provided by the case of peroxidase where the combination of hematin with the free protein is itself a slow reaction yet the compound is very stable. The great stability of the protein-iron protoporphyrin unit enables these hemoproteins t o form a wide variety of coordination complexes in straightforward bimolecular reactions involving replacement of the water molecule or OH group attached t o the sixth iron valency. The evidence that a water molecule or OH group is attached in this way is largely indirect (Keilin and Hartree, 112). The best evidence is in the case of hemoglobin. Haurowitz has shown t h a t when reduced hemoglobin is dried there is a very marked change in the absorption spectrum and the single diffuse band a t 555 mp is replaced by two sharp bands typical of a hemochromogen (4). Addition of water restores the reduced hemoglobin unchanged. Oxyhemoglobin does not show this behavior, its characteristic color and absorption spectrum are identicad in solution and in the dry state. Furthermore, dry oxyhemoglobin does not dissociate into hemoglobin and oxygen a t low oxygen pressures, but does if water is added. The essential oxygenation reaction is thus seen t o be: Globin.Fe,.(HnO)

+ + Globin.Fe,.Oz + H 0 2

2 0

where globin.Fe, represents one of the four heme groups of the hemoglobin molecule. Haurowitz obtained similar results with myoglobin and reconstituted hemoglobins containing protoheme, mesoheme, and the dimethyl ester of mesoheme. The hemochromogen formed when hemoglobin is dried was shown t o be a n intermolecular compound in

THE SPECIFIC REACTIONS O F IRON I N SOME HEMOPROTEINS

375

which a basic group of one molecule is joined t o the heme group of another, i.e., globin.Fe;globin.Fe. * . . , by the absence of any hemochromogen formation when glucose is present in the solution. It was t o be expected that the multiple polar groups of the glucose molecule would adsorb on the surface of the hemoglobin molecule while in solution and thus prevent the mutual association on drying. There is no similar evidence as yet for a water molecule attached t o the iron atom in methemoglobin but the change in color from brown t o red which occurs as a solution is made more alkaline is best explained by the ionization of such a water molecule. [Globin.FeP+H2O]S Globin.Fe,.OH

+ H+

The alkaline methemoglobin formed has a magnetic susceptibility of 4.45 Bohr magnetons compared with 5.77 for methemoglobin in moderately acid solutions. This marked difference is indicative t h a t the group ionizing is associated with the iron atom. Similar p K values of 8.10 a t 30°C. are obtained for the ionization by both spectroscopic and magnetic determinations (see Lemberg and Legge, 5 ) . I n the case of peroxidase the reduced form is assumed to have a water molecule attached, by analogy with hemoglobin. The oxidized form, according t o titration data obtained for the free protein and the reconstituted enzyme and a study of the dissociation of the fluoride complex, undergoes ionization of the water molecule with a p K of 5.0 (Theorell and Paul, 15; and Theorell, 115). There is no marked spectroscopic change, however, and this leaves unexplained a n ionization with a p K 10-11 which is accompanied by a change in color and magnetic properties. Lemberg and Legge have discussed these results fully ( 5 ) , and it seems more probable that the ionization of the water molecule corresponds t o this second pK. As was mentioned in the introduction, catalase is unique among the hemoproteins because its reduced form is unknown; all attempts t o prepare it from the free enzyme including electrolytic reduction have failed. On the other hand Keilin and Hartree (16,17) have found that in the presence of azide ions reduction does occur when hydrogen peroxide is added. The nature of this reaction is unknown. It may be a consequence of catalase possessing four hematin groups per molecule and that azide combined with one group alters the redox properties of the others in a manner analogous t o heme-heme interaction in the hemoglobin-oxygen reaction. This is unlikely since azide catalase like free catalase is resistant t o other reducing agents, even sodium hydrosulfite. An alternative explanation based on the azide ion undergoing oxidation in the reaction and t h e stabilization of the ferrous form of catalase by a

376

PHILIP GEORGE

reaction product of the azide ion has no experimental support, for no destruction of azide has been detected (Foulkes and Lemberg, 18). The question whether a water molecule is bound in this reduced form of catalase has clearly little meaning until the detailed chemistry of this reaction is known. Catalase itself has been shown by Agner and Theorell (19) t o have a hydroxyl group attached t o the iron protoporphyrini group with a pK of 3.8. The evidence for this rests on changes in light absorption on addition of different anions such as phosphate, acetate, and formate and the inhibition of catalase activity by these anions increasing as the hydroxyl ion concentration is decreased. Hemoglobin and myoglobin, peroxidase, and catalase combine with a very wide range of compounds in simple bimolecular reactions involving the replacement of the water molecule or OH grou:p. In general the ferrous compounds show preferential reaction with neutral molecules and the ferric compounds with anions, as can be seen for the typical compounds listed in Tables I1 and 111. The most ch:zracteristic properTABLE I1 Ferrous Compounds

Reagent

None 0 2

co

NO CNEtNC Ph.NO

Hemoglobin

m.i. s.d., m.c. s.d., m.c. s.d., m.c. s.d., m.c. s.d., m.c. s.d.

Reduced Peroxidase

m.i.

-

s.d., m.c.

Spectroscopically similar to parent compound. 8.8. s.d. Spectroscopically different to parent compound. m i . or m.c. Bonds to iron atom shown to be essentially ionic or ,covalent by magnetic ausceptibility measurements. A dash indicates that the compound is not formed, a blank space that it has not been investigated.

Note.

ties of these compounds are their color, type of absorpt8ionspectrum, and magnetic susceptibility. The six bonds around th.e iron atom are predominantly ionic in the parent compounds as shown by the susceptibility values which correspond very closely to th.e expected values for four and five unpaired electrons in the ferrous and ferric parent compounds respectively. Those compounds in which the bonding is still ionic have absorption spectra and colors similar t o the parent compounds. The formation of many compounds, on the other hand, results

THE SPECIFIC REACTIONS OF IRON I N SOME HEMOPROTEINS

377

TABLE I11 Ferric Compounds Reagent

Methemoglobin

Peroxidase

Catalase

None OHAnions FCNHSN 1NHzOH N2H4 NO NH 8 EtOH Imidarole Cyanate Thiocyanate

m.i. s.d., m i . &

m.i. s.d.

m.i. parent cpd.

s.s., m.1. s.d., m.c. s.d., m.c. s.d., m.c.

s.s., m.i. s.d., m.c. s.d., m.c.

s.s., m.i. s.d. s.d., m.c. s.s., m.1.

9.8.

S.S.

9.9.

8.5.

9.8.

s.d. m.c. m.i. s.d., m.c.

s.d.

s.d. 9.9.

S.8.

S.S.

Note. 8.8. Spectroscopically similar to parent compound. s.d. Spectroscopically different t o parent compound.

m i . or m.c. Bonds to iron atom shown t o be essentially ionic or covalent by magnetic susceptibility measurements. a Alkaline methemoglobin has an anomalous susceptibility of 4.47 Bohr magnetons corresponding most closely to the value for three unpaired electrons (see Pauling, 4).

in a large decrease in magnetic susceptibility corresponding to no unpaired electrons in the ferrous compounds and one unpaired electron in the ferric compounds. These are the values expected for covalent octahedral iron complexes. This susceptibility change is usually accompanied by a profound alteration in color, type of absorption spectrum, and the intensity of the absorption bands which is usually enhanced. Tables I1 and 111indicate those compounds which are spectroscopically similar (s.s.) and those which are spectroscopically different (s.d.) from the parent compounds, and in addition whether magnetic measurements have shown the compounds to contain essentially ionic (m.i.) or covalent (m.c.) bonds around the iron atom. Table IV which is based on a general classification suggested by Theorell (20) and Hartree (21) shows the predominant correlation between color and spectrum type. I n addition t o the bands in the visible region of the spectrum all these compounds have a very intense band between 380 and 440 mp, in the so-called Soret region, where position and intensity is characteristic for each compound. I n some spectroscopic analyses it has proved more satisfactory t o work at these wavelengths. Hartree has recently reviewed the magnetic properties of hematin derivatives (21) and a detailed

378

PHILIP GEORGE

TABLE I V General Correlation between Bond T y p e , Color, and Absorption Spectrum (Weaker absorption bands are indicated by brackets; the wavelengths are given in millimicrons) Color

Spectrum Type

Examples with the Band Maxima

Ionic ferric compounds

Green-brown

Absorption band in MetHb; 637, (582), (548), red between 600 504 and 640 mp; strong Catalase: 639, (544), 506 band in blue some- Fluoride catalase: (622), 597, times faint bands (542) in the green Aside ca.talase: 624, (581), (567), (.536), 494 Peroxidme: 645, 583, 548, 498

Covalent ferric compounds

Bright red

Two diffuse bands in the green

Aside MetHb: 575, 542 HS. MetHb: 570, 545 Cyanide peroxidase: 581.5, 542

Ionic ferrous compounds

Carmine-red to purple

Diffuse band in the green

Hb: 555

Covalent ferrous compounds

Scarlet to pink

Two very sharp bands in the green

HbOz: 5‘75, 540 HbCO: 670, 540 HbNO: 568, 531 CO peroxidase: 578, 545.5

Note. Alkaline methemoglobin has the spectrum typical of a covalent ferric eornpoiind y e t its magnetic susceptibility value is anomalous (Table 111).

description of all the various compounds with references t o the original investigations is given by Lemberg and Legge ( 5 ) . Mention has been made above that the formation and properties of these compounds show them t o be octahedral coordination complexes in which the bonds t o the central iron atom are either essentially ionic or essentially covalent. There is, however, no correlaiion between the type of bond being formed in the complex and the speed of the reaction. The kinetic study of the reaction between hemoglobin and oxygen or carbon monoxide prompted the development of the rapid flow ” technique by Hartridge, Roughton, and Millikan for following fast reactions in solution (see Roughton, 4). These kinetic investigations showed both combination and dissociation reaction t o be very rapid. A comparison of velocity constants for hemoglobin and myoglobin reacting with O2 and CO is given in Table V based on data compiled by Millikan (22). ((

THE SPECIFIC REACTIONS OF IRON I N SOME HEMOPROTEINS

379

TABLE V Velocity Constants for Hemoglobin and Myoglobin Reactions at 2O"C., p H 7.4 Velocity Constants lteaction

O2 combination O p dissociation CO combination CO dissociation

1.9

3 4

Mb

Hb

x

106 M.-1 set.-' 40 sec.-l 1.3 X lo5 M.-1 set.-' 4 X sec.-l

107

37

x 106 x 10-2

4

x

The combination reactions involve a change in bond type from ionic to covalent and the dissociation reaction the reverse : Hb(H20) (Ionic)

+

0 2

HbOz

+ Hz0

(Covalent)

but it is clear from the data in Table V that both reactions are rapid. Eley (23) has examined the kinetic data on these reactions and has shown that the combination reactions are fast on account of low heats of activation and the dissociation reactions fast because of fairly high positive entropies of activation. It is thus not possible t o deduce from the speed of one of these reactions the change in bond type occurring. Apart from these reactions of hemoglobin and myoglobin, and the formation and subsequent reactions of the hydrogen peroxide complexes of peroxidase and catalase which have been intensively studied by Chance and will be described later, few have been examined kinetically. There are some kinetic studies for instance on the reduction of methemoglobin by hydrosulfite, the oxidation of hemoglobin by ferricyanide ions and the formation of methemoglobin cyanide. However, there are qualitative indications that the formation and dissociation of the complexes listed in Tables I1 and I11 are all fast reactions. Hence when a reaction is found t o be slow and the dissociation of some of the peroxide complexes of peroxidase and catalase are slow, it suggests that the complex may not be formed in a simple bimolecular displacement reaction like the formation of HbOz or HbCO. The rapidity of the hemoglobin reaction is a fair indication that the heme groups lie flat on the surface of the protein molecule. The observations of Perutz on the behavior of crystalline acid methaemoglobin offer further support (4). X-ray studies show that hemoglobin molecules are rigid and impenetrable t o liquid and crystalline oxy- or carbonmonoxyhemoglobin can be transformed into methemoglobin without any change in crystal structure. Acid methemoglobin crystals suspended

380

PHILIP GEORGE

in ammonium sulfate solution can be converted into alkaline methemoglobin by the addition of diammonium phosphate, and into azide methemoglobin by the addition of sodium azide without causing any change in unit cell dimensions. Thus these reactions between hemin groups and ions diffusing through the liquid of crystallization show the hemin groups to be located on the surface, and the constant unit cell dimensions show that these groups do not play a part in the bonding between neighboring hemoglobin molecules in the crystal lattice. In contrast t o the simplicity of these results it should be mentioned that in some way the crystal structure of reduced hemoglobin is different. Attempts to convert crystals of acid methemoglobin into reduced hemoglobin and crystals of reduced hemoglobin into oxyhemoglobin are always accompanied by a complete break up of the original crystal structure and recrystallization in a new form occurEi (Haurowitz, 24). It is not yet established whether these differences are only important in considering the structures of the crystals or whether they are reflected in the reactions of these compounds in solution. I n studying the reaction of the heme group in hemoproteins it is always necessary to ensure that no alteration in protein structure is affecting the results. Complete denaturation is easily recognized by precipitation and flocculation and the liberation of heme of hematin, but minor changes can also occur. I n the case of peroxidase two forms were isolated by Theorell (25) and named peroxidase I and 11. The first of these, now called paraperoxidase (Theorell, ZS), has been found t o have undergone some alteration possibly the removal of the carbohydrate portion of the protein molecule or possibly reversible denaturation t o a slight extent (Keilin and Hartree, 48). The paraperoxidase shows enzymatic activity comparable to that of the intact enzyme. Oxyhemoglobin can also undergo minor changes in structure which are apparent in a slight shift of the Soret band maximum absorption first from 414.5 mp t o about 410 mp, and then t o 406 mp (Jope, 4). It is an interesting feature that hemoglobin reconstituted from heme and free globin also shows maximum absorption at 410 mp. In addition to structural changes of this kind occurring, when the hemoproteins take part in oxidation-reduction react ions there is the further possibility of irreversible oxidative attack on the porphyrin ring itself. Lemberg, Legge, and Lockwood (27) have studied this process in the case of hemoglobin and have named the oxidation product choleglobin: the precise nature of this compound has not yet been established. The reactions described above and the specific reactions described below are to be regarded as characteristic of the four hemoproteins i n general. There are often differences in amino acid content and chemical

THE SPECIFIC REACTIONS

OF IRON IN SOME HEMOPROTEINS

381

behavior of the ‘ I same ” hemoprotein derived from different sources. For instance there are wide variations in the oxygen dissociation curves of mammalian and avian hemoglobins. Peroxidase obtained from horseradish has different properties from leucocyte peroxidase which contains a different porphyrin group (Lemberg and Legge, 5 ) . Bacterial catalase has a little greater catalatic activity than erythrocyte catalase although both hemoproteins contain four intact hematin groups; and both are more active than liver catalase in which one or more of the hematin groups have been converted into a catalytically inactive bile pigment molecule. These differences are described in detail by Lemberg and Legge (5). The reactions described in this paper refer mainly to horse hemoglobin and myoglobin, horseradish peroxidase, and catalase obtained from horse liver or erythrocytes. AND MYOGLOBIN AUTOXIDATION AND OTHER 111. HEMOGLOBIN REACTIONS

The reversible combination of these hemoproteins with oxygen and carbon monoxide has been very extensively studied. Reviews of kinetics and equilibrium measurements are found in Haemoglobin (4), Lemberg and Legge ( 5 ) , and Wyman (1). The main themes of these studies have been t o explain the sigmoid shape of the oxygen equilibrium curve of hemoglobin which is in marked contrast with the normal hyperbolic curve found for myoglobin and t o explain the change in acidity which accompanies the oxygenation of hemoglobin, both of which are very important aspects of its biological function as an oxygen carrier. In hemoglobin it has been shown that interaction occurs between the heme groups such that when one has combined with oxygen the affinity of the remaining groups is altered. Several estimates have been made of the equilibrium constants for the formation of the successive intermediates Hb4(02),Hb4(02)2,Hbr(02)% and Hb4(0& (see Roughton, 4). Similar interaction occurs in the redox system hemoglobin-methemoglobin (Wyman, 1). There is a possibility of a similar sort of interaction occurring in catalase which has four hematin groups per molecule, and mention will be made of this later. In the case of hemoglobin it has been suggested that some oxidation-reduction reactions can proceed by an intramolecular mechanism on account of four reacting heme groups being attached to the same molecule. This mechanism has been developed particularly to account for the autoxidation of hemoglobin by Lemberg and Legge (5). The autoxidation reaction has received relatively little attention, but since it has considerable bearing on the problem of why oxyhemoglobin is stable, it will now be discussed in detail. It is often stated that the globin protects the heme from oxidation,

382

PHILIP GEORGE

and t h e problem is usually regarded as structural, the nature of the bond linking the ferrous protoporphyrin group to protein, conferring the highly specialized activity on the central iron ato:m. The way this happens may be considered as follows. The existence of oxyhemoglobin as a stable molecule is a particular case of the general problem of the stability of coordination compounds. As is well-known, this stability is very dependent on the nature of the coordinating groups and though combination with molecular oxygen is rare the ability of hemoglobin t o combine with it can reasonably be attributed to a favorable balance of bond energies, including as the crucial one that of the heme-protein link. The resistance to oxidation is also in part a structural problem for the nature of the coordinating group around II metal ion affect its oxidationreduction potential which is in part a function of the more fundamental property of the ion, the ionization potential of its reduced form in solution. The magnitude of this ionization potential immediately determines whether the first step in the oxidation is exothermic and so possibly very rapid, or how endothermic it is which would give an indication of how slow the step might be. But apart from this structural aspect of the resistance to oxidation there is also a kinetic aspect. Hemoglobin in the presence of oxygen does form methemoglobin, and so the mechanism of the oxidation, revealed by an analysis of the kinetics, should provide additional data for understanding the slowness of the oxidation when compared with other oxidation processes. Brooks (28,29) showed that the oxidation of hemoglobin at pH 5.69 is first order in unoxidized hemoglobin and the first order velocity conatant kob.. is dependent on the oxygen pressure according t o the complex function illustrated in Fig. 1A. Brooks summarized his results in the following rate equation:

where a is the fraction of uncombined hemoglobin, (a - t) the percentage of unoxidixed hemoglobin at time t , p the oxygen pressure, and k' and b numerical constants. In addition he made observations on the effect of varying the hydrogen ion concentration-the reaction is faster in more acid solution-the concentration of buffer salts, the effect of additions of sodium chloride, and the effect of temperature. Coryell, Stitt, and F'auling (30) later confirmed the first order dependence on unoxidized hemoglobin. Conant and Fieser (31) considered three possible paths the reaction could take: (a) the spontaneous decomposition of Hb02 into methemoglobin, ( b ) the oxidation of Hb by HbO2, ( c ) the direct oxidation of Hb

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

383

by 0 2 . In the light of his experimental results Brooks favored a mechanism based on ( c ) , for in (a) kObs.would increase t o a maximum value with increasing oxygen pressure, and in ( b ) the reaction would be second order in unoxidized hemoglobin. Direct reaction between Hb and 0 2 accounted for the term a(a - z ) p in his rate equation and left the b / ( l bp) term t o be explained. Brooks explored the possibility of certain intermediates in t,he oxygenation sequence Hb4(02)I, Hb4(02)2, Hb4(02)3 reacting preferentially with 02,and showed that no such selective reaction could account for the results. He put forward two possible explanations for the b / ( l b p ) term: (i) The presence of an oxidation catalyst R, in equilibrium with RO2 which oxidizes Hb (cf. Neil1 and Hastings, 32). (ii) Some type of chain mechanism for the oxidation in which O2 acts as an inhibitor. No further evidence has been found for an independent oxidation catalyst and recent work on autoxidation sheds no light on a chain mechanism which would explain the results. Lemberg and Legge ( 5 ) proposed that the oxidation catalyst is HbOz itself, but this point will be discussed later. I n 1942 Legge (33) suggested the spontaneous decomposition of one of the intermediates in the oxygenation sequence into methemoglobin as a possible mechanism. Using Pauling’s equation (34) for the fraction of hemoglobin in the form Hb4(02)1, Hb4(02)2, etc., as a function of oxygen pressure, he showed that the spontaneous breakage of Hb4(02)2 gave good agreement with Brooks’ data at low oxygen pressures. Lemberg and Legge ( 5 ) later amplified this mechanism by picturing the Fe2+02groups as oxidative catalysts and accounted for the selective action of Hb4(02)2by assuming the participation of two oxidizable XH2 groups on the protein t o make an intramolecular reaction possible fitting the stoichiometric relation:

+

+

+ 2XH2

Hbd(Oz)P

-+

Hb4(0H)r

+ 2X

in which the essential valency change occurring is: Fe2+(02).

. . XH2 . .

+

Fe2+--, Fe3+(0H) .

*

*

X

. . . Fea+(OH)

where Fez+, Fe2+02,and Fe3+(OH) represent the protoporphyrin iron atoms in hemoglobin, oxyhemoglobin, and methemoglobin respectively. Two particular assumptions underlie an intramolecular mechanism of this kind: (a) Fe2+and Fe2+02groups can react with each other (and with XH2 in the later mechanism) even though they are fixed on the globin molecule at considerable distances from each other. ( b ) I n spite of many collisions with other hemoglobin molecules, no significant inter-

384

PHILIP GEORGE

molecular reaction occurs. This intramolecular change is a necessary assumption in these cases to account for the first order dependence in unoxidized hemoglobin. The existence of this type of reaction is not yet established, and one difficulty in accepting it is the necessary assumption ( b ) that a normal reaction path by collision is blaicked in some way. On kinetic considerations alone the hypothesis of XH2 groups participating in the reaction is unnecessary. A mechanism of this kind has better claims if H 2 0 2is the reaction product and XH2 groups play no part at all, because a reaction path involving such groups would be an addition t o potential reaction paths similar to those responsible for the autoxidation of heme itself or hemochromogens which are rapid reactions. Bringing XH2 groups into the reaction mechanism makes it more difficult to understand why the reaction is slow. The validity of this mechanism rests primarily on its ability to explain the experimental results and over the low oxygen pressure range the agreement is good. However, if similar calculations are made for the higher pressure data serious discrepancies are found, especially at 723 mm. 0 2 where the reaction proceeds about forty times faster than it should according t o Hb4(02)2 decomposing (Brooks, 35). Even though the Pauling equation used in these calculations must be abandoned in the light of recent equilibrium measurements (Roughton, 4), there is no doubt that the fraction of hemoglobin present as Hb4(02)2 at high oxygen pressures will still be extremely small and this discrepancy in the oxidation rate still .appear. It seems doubtful therefore whether Hb4(02)2does play this unique role. An alternative approach to the problem was suggestNedby George (36) who showed that Brooks observed first order constant icob.. has an oxygen pressure variation very close t o that given by the product of the concentration of Hb and HbO2, i.e., a(1 - a). The maximum value of kob..should thus occur at the 0 2 pressure for half saturation in accord with experiment as in Fig. 1A. However a kinetic paradox appears for a velocity constant proportional t o [Hb] X [Hb02] should result from a reaction that is second order in unoxidized hemoglobin and not first order as was fully established by Brooks. This paradox may be resolved if the rate equation is written:

and a chemical mechanism looked for which would lead t o the concentration of unoxidized hemoglobin appearing in the denominator. Such a mechanism will be discussed later. George and Stratmann (37) investigating the kinetics of the oxidation of myoglobin t o metmyoglobin by molecular oxygen, under the aame

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

385

100

0 OXYGEN PRESSURE, mm.

ij i

A: HEMOGLOBIN

l I

I I

0

I

I

I

I

FIG. 1. Variation of the first order rate constant with oxygen pressure for: A, hemoglobin (Brooks, 29) ; B, myoglobin (George and Stratmann, 37). Temperature 30"C.,pH = 5.69. The broken lines indicate percentage saturation determined a t the various oxygen pressures. For rnyoglobin the two lines show the experimental variation.

conditions that Brooks used, found similar behavior. The reaction which proceeds about six times faster is first order in unoxidized myoglobin and the first order constant shows the same type of variation with oxygen pressure (see Fig. 1A and B). The maximum rate again occurs at the oxygen pressure required for half saturation indicating that both Fe022+ and free Fe2+ groups are involved in the reaction. Since myo-

386

PHILIP GEORGE

globin contains only one heme group on each protein imolecule an intramolecular mechanism such as t h a t suggested by Lemberg and Legge ( 5 ) is ruled out and for myoglobin the mechanism must be intermolecular. If it is accepted that similar oxidation kinetics result from an identical reaction mechanism, and the more complicated the kinetics the more likely this becomes, then it is very probable that with hemoglobin too t h e reaction is intermolecular. A system which shows some resemblance t o the autoxidation reaction is the oxidation of various substrates in the presence of hemoglobin and oxygen accompanied by the transformation of hemoglobin into choleglobin. The coupled oxidation of ascorbic acid in this system has been extensively investigated by Lemberg, Legge, and Lockwood (27), who compared its action with the production of verdohemochrome when hydrogen peroxide is added t o hemochromogens in the presence of hydrogen donors such as ascorbic acid (Lemberg, Cortis-Jones, and Norrie, 14). They showed t h a t the coupled oxidation is initiated by t,he direct reaction of oxyhemoglobin with ascorbic acid and suggested t h a t the reaction could be formulated:

+ tFez+ + A < Fez+

0 2

HrA

~

[Fez+02] HrA

A

+ [Fe2+H2021-+ FeS+

Cho:leglobin

*(

The ratio of ascorbic acid oxidized t o choleglobin formed is about 10: 1. B y using methemoglobin in place of oxyhemoglobin they found t h a t the rate of choleglobin formation was the same. Since Holden had established that the coupled oxidation could be poisoned by cyanide (38) this is evidence that methemoglobin acts as a true reaction intermediate. No detailed kinetic study of the system has been carried out but Lemberg, Legge, and Lockwood (39) showed that a t an oxygen pressure of 15 mm. the reaction proceeded about four times faster than at 150 mm. which indicates that the reaction mechanism is rather inore complicated than the above scheme where the reaction rate should rise t o a maximum with increasing oxygen pressure. The similarity with the autoxidation kinetics suggests that both H b and HbOz play a part in the reaction. A far more rapid oxidation of ascorbic acid occurs when it is mixed with a solution of oxyhemoglobin and acid then added t o denature the protein. I n the absence of ascorbic acid, oxyhemoglobin yields denatured protein and acid hematin, only 40% of the oxygen is evolved, hydrogen donor groups in the globin are oxidized, and a small part of the hematin also undergoes oxidation with liberation of its iron. If ascorbic acid is present it is oxidized instead of the globin. The oxidation occurs instantaneously if the ascorbic acid is added before acidification, but much

THE SPECIFIC

REACTIONS

OF IRON IN SOME HEMOPROTEINS

387

less is oxidized if it is added after. Lemberg and Legge ( 5 ) suggest that the ascorbic acid reacts directly with activated oxygen in an oxyhemoglobin-ascorbic acid complex, the oxygen being activated as a result of the acid altering the oxygen-heme link. There is a second way in which this system can oxidize substrates. With biliverdin and bilirubin (open-chain derivatives of porphyrins) there is no immediate decay of the oxidizing system after acidification. Lemberg (40) and Lemberg, Legge, and Lockwood (41) showed that this was due t o the hydrogen peroxide present produced by the action of acid on oxyhemoglobin, oxidizing the substances with the acid hematin also produced acting as a peroxidase. A possible free radical interpretation of these reactions will be discussed later. A different type of reaction is that between methemoglobin and hydrogen peroxide. Kobert (42) showed that a red-colored complex was formed and Keilin and Hartree (43) demonstrated that one peroxide molecule per iron atom was required. The complex which has absorption maxima at 585 mp (a band) and 545 mp (0band) slowly decomposes, liberating methemoglobin, but is rapidly destroyed by hydrogen donors such as ascorbic acid and hydroquinone. Haurowitz (44) claimed that addition of alkali t o the complex yielded alkaline methemoglobin and represented the reaction:

+ OH-

MetHb.+HzOz

4

MetHb.OH

+ Hz02

Keilin and Hartree showed too that ethyl hydrogen peroxide gives a complex with absorption maxima a t the same wavelengths. I n a more recent paper (45)) they have extended their observation on the hydrogen peroxide complex, confirming their results at low HzOz/MetHb concentration ratios. As this ratio is increased they found that the absorption spectrum is modified, the a band which previously showed a maximum at 585 mp splitting into one with two maxima at 578 and 592 mp. Increasing the HzOz/MetHb concentration ratio still further, resulted in the disappearance of the longer wavelength band and the spectrum with maxima now at 578 mp (aband) and 545 mp ( p band) was recognized as that of oxyhemoglobin. At the concentration of hydrogen peroxide required to obtain the oxyhemoglobin spectrum catalytic decomposition of the peroxide occurred. This establishes beyond any doubt that the catalytic decomposition of peroxide in this case proceeds by a mechanism involving a valency change of the iron atom. Keilin and Hartree (45) showed too that metmyoglobin forms a similar compound with hydrogen peroxide which undergoes the same reactions. George and Irvine (116) found that this complex is formed at all hydrogen ion concentrations between pH 5.5 and 12.0, but that only

388

PHILIP QEORQE

when the formation occurs in the pH range 8-9 can complete regeneration of the metmyoglobin be effected by adding mild reducing agents. The spectrum of the complex in this region shows a sharp band at 549 mp with emM. = 9.8 and a broad band appearing as a shoulder at 570-585 mp with cmy. = 8.4. In the Soret region there is a sharp band at 423 mp with emy. = 107. With solutions of pH < 8 some complex is formed, but an additional reaction involving the hematin also occurs, giving an equally intense spectrum but with the bands shifted. When the reaction is carried out in solutions of increasing alkalinity above pH 9.0 progressively less complex is formed and a hematin degradation product, colorless in both the visible and Soret regions of the spectrum, is produced. Titration of the complex using ferrocyanide as the reducing agent showed that one ferrocyanide ion reacts per molecule of the complex. The complex thus has only one oxidizing equivalent compared with the two oxidizing equivalents of hydrogen peroxide itself. On the other hand, when varying amounts of ferrocyanide and metmyoglobin were mixed and hydrogen peroxide then added, the full oxidizing capacity of the peroxide was obtained. This suggests very strongly that the formation of the complex proceeds in a reaction of the type MetMb

+ HzOz

(2 oxid. equiv.)

--*

+

Complex X (1 oxid. equiv.) (I oxid. equiv.)

and that a transient oxidizing entity X is produced along with the complex, which can react with added reducing agents if present initially, but which disappears if the complex formation is allowed t o proceed independently. These results indicate that the complex cannot be of the ion pair type [MetMb-OOH] for such a compound would have two oxidizing equivalents. A tentative mechanism which can account for these results involving the formation of a quadrivalent iron compound “ferrylmyoglobin ” will be discussed later.

REACTIONS IV. PEROXIDASE The catalytic activity of peroxidase is intimately connected with its ability to form complexes with hydrogen peroxide. Three such complexes are formed, depending on the experimental conditions and they are interrelated :

+

Peroxidase Ha09 + Per. (HnOi) I (brown) (1 mole) green

excem

-

Per. (Hz01) 11 Per. (HaOa) I11 H:Oi pale red deep red

spontaneous

decomposition

Per. (HzOz) I1 pale red

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

389

As the structure of these complexes is unknown they are denoted as I, 11, and 111. Keilin and Mann in 1935 described complexes I1 and I11 (46). Complex I1 is formed from equimolar amounts of peroxidase and hydrogen peroxide and with an excess of peroxide it changes into complex 111. Both these complexes react rapidly with hydrogen donors. Carbon monoxide does not affect their spectrum, which led Keilin and Mann (46) to suggest that the iron atom is in the trivalent state. Later in 1942 Theorell (20,47) observed that the pale red complex I1 was not formed immediately when peroxidase was added to hydrogen peroxide but that a transient green-colored complex was formed first which very rapidly changed into the pale red complex 11. The position of the absorption maxima of these compounds together with those of free peroxidase in the visible spectrum are listed in Table VI. TABLE VI Compound of Peroxidase Neutral peroxidase Peroxidase (H20z)I Peroxidase (H202)I1 Peroxidase (H2Oe) I11

.

Position of Absorption Maxima (mI*)

Reference

645, 583, 548, 498 658 - - - 561, 530.5 - 555," 527 - 583, 545.5 583, 546

46 20,47 46 48 46 48

a Asymmetric absorption band, 555 mp, refera to the observed maximum, not the "center" of the band.

Chance (49,50) has extended the spectroscopic examination of these complexes to the Soret region 370-450 mp and has examined too the complexes with methyl and ethyl hydroperoxide. The primary complexes are also green and have nearly identical spectra in the Soret region: these maxima are at 407 mp. The Soret band of the free enzyme is considerably diminished in intensity on forming the complex, and it is slightly shifted toward the visible. The alkyl hydroperoxides also yield pale red secondary complexes with spectra in the visible region very similar to the hydrogen peroxide complex 11. In the Soret region all have about the same intensity as the free enzyme, but there is a shift of some 15 mp toward the visible. Chance comments that the identity of the Soret bands of the peroxide complexes with HzOz and methyl or ethyl hydrogen peroxide indicates that the nature of the ironperoxide bond is unaffected by substitution in the Hz02 molecule. Theorell (20,51) has suggested on the basis of a magnetochemical study that the bonding in the pale red secondary complexes is covalent, whereas

390

PHILIP GEORGE

a comparison of the spectra of the green primary complexes with the green peroxidase-fluoride complex, which is known t o contain ionic bonds, indicates that the bonding in these primary complexes is ionic. Chance (50) found that the alkyl hydroperoxides do not give a deep red-colored complex analogous t o peroxidase-hydrogen peroxide 111. Instead a bright green complex with a very strong band at 675 mp and a faint band at 557 mp is formed. This complex which he designates complex IV has a Soret band resembling that of the primary complexes. Since there is considerable destruction of peroxidase at th e large peroxidase concentrations required t o form this complex, hLe suggests th a t it is related t o t he green oxidation products of hemoglobin and catalase studied by Lemberg and his associates (see Lemberg and Legge, 5). Abrams, Altschul, and Hogness (52) suggested th a t the similarity between the visible bands in the pale red complex given by the peroxidase which can oxidize cytochrome c, and th e bands in oxymyoglobin and oxyhemoglobin indicates th at the secondary complexes may be ferrous compounds in spite of their resistance t o carbon monoxide. However, Chance points out that this spectroscopic analogy does not hold in the Soret region 370-450 mp. I n this region there is some resemblance when the bands of metmyoglobin in relation t o oxymyoglobin are compared with those of peroxidase in relation t o the primary complex. These inconsistencies in spectral analogies when th e visible region is compared with the Soret region thus makes it impossible t o draw any definite conclusions. Some of the early studies of the kinetics of peroxidase oxidations provided evidence t o suggest that both the hydrogen peroxide and the hydrogen donor were joined to the enzyme molecule (Mann, 53). It was found that an excess of hydrogen peroxide inhibited the oxidation of the hydrogen donor, and th at this inhibition could be relieved by increasing the concentration of the hydrogen donor. This could be explained if hydrogen peroxide first combined with the enzyme and then peroxide and the hydrogen donor competed for a second site on th e enzyme. Combination between the hydrogen donor and the enzyme has been accepted by Lemberg and Legge (5) in their theory of peroxidase action which they represent: Fe3+OH

FeJfOH HzA

Fe3+OOH HzA

FeSfO H

I

+ HzO + A

However, the inhibition of peroxidase action a t high peroxide concentrations is capable of other interpretations, and since there is no evidence for the enzyme uniting with t.he hydrogen donor throiigh a group in its

THE SPECIFIC REACTIONS

OF IRON IN SOME HEMOPROTEINS

391

protein in the detailed kinetic studies of Chance (see below) the acceptance of this mechanism must await a more conclusive proof of such combination. Chance has recently extended his original observations on the kinetics of the formation of the pale red peroxidase-H202 complex and its reactions with hydrogen donors like leucomalachite green and ascorbic acid (54) t o include the kinetics of formation of the green primary complex and its conversion into the pale red secondary complex (55) in the case of both hydrogen peroxide and alkyl hydroperoxides. Using a rapidflow apparatus equipped with a monochromator, kinetic data were obtained in agreement with the following basic reaction mechanism : Per. OH

+ HOOR kikz Per. OOR (I) + H2O kl

Per. OOR (I) + Per. OOR (11) ka

Per. OOR (11) -+ Per. (OH) Per. OOR (11)

+ AH2-+ Per. (OH) + ROH + A ks'

(2)

(3) (3')

The secondary complex decomposes spontaneously as in 3 and can react with hydrogen donors as in 3'. The study of the last reaction entailed using concentrations of hydrogen donors (AH2) in excess of t h a t of the secondary complex such that the kinetics were pseudo first order. The net velocity constant k,' is thus equal t o kd(AH2) where kk is the true bimolecular constant for this reaction. The values of kl were found t o be 9 k 2 X lo6, 1.5 X lo6 and 3.6 X 106 M.-l sec.-l for hydrogen peroxide and methyl and ethyl hydrop.eroxides respectively; and k , the first order constant for the conversion of the primary into the secondary complex was about 4 sec.-l for all three peroxides. The spontaneous decomposition of the secondary complex depends principally upon whether peroxides have been previously added t o the solution. A fresh peroxidase gives velocity constants for k , of about 0.02 set.-', but several additions of peroxide can reduce this velocity constant by a factor of ten. T o explain this effect Chance suggests t h a t a hydrogen donor is present with the enzyme initially which can accentuate the decomposition of the secondary complex as in reaction 3', but with successive additions of peroxide it becomes oxidized. This basic mechanism holds if the reaction is carried out in two stages: first, the formation of the primary green complex followed by its quantitative conversion into the secondary complex, then addition of the hydrogen donor which is oxidized by the secondary complex. However, a more complicated mechanism is possible for Chance also found that the conversion of the primary complex into the red secondary complex is accentuated in the presence of a

392

PHILIP GEORGE

hydrogen donor, and the velocity of formation of the secondary complex from the primary complex can be nearly as rapid as the actual formation of the primary complex. The value of kz, the velocity constant for the decomposition of the primary complex into peroxidase and peroxide cannot be obtained by direct measurement. Measurements of the equilibrium constant for the dissociation of the primary and secondary complexes with methyl M. and 3 X lo-' M. from which hydroperoxide gave values of 3.2 X kz may be estimated to be 2.2 sec.-l or 3.4 sec.-' from the known formation velocity constants. Chance found that pH has little effect on the dissociation constants in the range 3.6 to 8.8, which suggests that neither of the two hemelinked groups found by Theorell and Paul (15) affects the ability of peroxidase to form its enzyme substrate complexes. On the basis of this observation he suggests that the primary combination takes the form: Per. OH

+ HOOR kikl Per. OOR (I) f H1O

Above pH 8.8, where further decrease of hydrogen ion concentration leads to increased formation of alkaline peroxidase, peroxide complex formation is suppressed, which indicates that peroxides cannot replace the covalently bound hydroxyl group. There is an interesting parallel here in the case of methemoglobin where Haurowitz claimed that the alkaline form cannot give a peroxide complex (44). The reactions of the secondary complexes with hydrogen donors are rapid bimolecular reactions. Chance (49) has obtained the velocity constants for horseradish peroxidase, hydrogen peroxide and methyl and ethyl hydroperoxide complexes respectively reacting with ascorbic acid at pH 7.0: k4 = 2.8 X lo3, 2.8 X lo3, and 2.2 X lo3 M.-l sec.-l and reacting with pyrogallol at pH 7.0, k4 = 2.1 X IOs, 2.1 X lo6, and 1.8 X lo6 M.-l set.-'. With the three hydrogen donors, hydroquinone, guaiacol, and pyrogallol, there was no systematic decrease in the values of kb over the pH range 3.6-6.7. Whereas the peroxidase reaction is reversibly and competitively inhibited by cyanide ions (Lemberg and Legge, 5 ; Keilin and Mann, 46; Chance, 56) carbon monoxide has no inhibiting effect (Theorell, 47), which has led to the widely accepted view that the enzyme remains in the ferric state throughout the reaction. A very different type of reaction that peroxidase can catalyze is the autoxidation of dihydroxymaleic acid HOOC.C(O'K)=C(OH)-COOH which has the same oxidizable group as ascorbic acid (see Lemberg and Legge, 5). I n contrast to the usual peroxidase reaction with peroxides,

THE SPECIFIC REACTIONS O F IRON I N SOME HEMOPROTEINS

393

this system can be poisoned with carbon monoxide, and this inhibition is reversed by light. It would seem therefore that a valency change occurs in the course of this reaction. Theorell and Swedin (117) have shown that the brown color of free peroxidase changes to deep red when oxygen is bubbled through a mixture of peroxidase and dihydroxymaleic acid, and that the absorption bands of the red substance correspond to those of peroxidase in the presence of an excess of hydrogen peroxide, i.e., the third comRlex, Per. (HzOz)111. The initiation of the polymerization of vinyl compounds is a good indication of the production of OH radicals (Dainton, 108) particularly in systems containing hydrogen peroxide and ferrous ion (Baxendale, Evans, and Park, 84). It is interesting therefore that peroxidase and catalase with hydrogen peroxide have been found inactive (Dainton and Smith, 107). Another difference between the reactions of the enzymes and ionic ion, where the mechanism is generally accepted t o involve free radicals, appears when the products of hydrogen peroxide oxidation obtained using peroxidase are compared with those obtained using ferrous iron. Mann and Saunders (log), Saunders and Mann (110), and Chapman and Saunders (111) investigated the oxidation of aniline, p-toluidine, and mesidine by these two systems and found very different reaction products in the two cases. V. CATALASE REACTIONS

I . Catalytic Decomposition of Hydrogen Peroxide Many kinetic studies have been made of this reaction which is complicated by the destruction of the enzyme as the reaction proceeds. Typical of the early investigations are those by Yamasaki (57), Morgulis (58), Northrop (59), Williams (60), Nosaka (61), Maximovitsch and Antonomova (62), and Zeile and Hellstrom (63). Under certain experimental conditions the following kinetic equations hold: and

-d(H202)/dt = k[Catalase] . [H~OZ]

-d(Catalase)/dt

=

k’[Catalase] . [ H z ~ z ]

I

I1

The enzyme destruction particularly at a temperature of 20°C. or below is a relatively slow reaction compared with the peroxide decomposition, and since all the evidence shows it t o be a true side reaction, it will not be discussed further here. A change in activity of the enzyme observable at high peroxide concentrations but distinct from the slow destruction reaction will be referred to Iater. Measurements of peroxide decomposition under the conditions where Eq. (I) holds have been widely used for estimation of catalase activity

394

PHILIP QEORGE

according to the procedure suggested by Zeile and Hellstrom (63). This entails using very dilute catalase solution about 5 X 10-l1 M , and N HzO2 and estimating the residual peroxide at approximately three minute intervals after the reaction is started. This procedure has recently been criticized by Bonnichsen, Chance, and Theorell (64), who showed that Zeile and Hellstrom's procedure gave values for the enzyme activity about 60% of that obtained when higher enzyme concentrations were used, e.g., 2 X lou9 M , and samples withdrawn. for titration at 13, 28, and 43 seconds from the start of the reaction. 'They suggest that enzyme inactivation arising from the spreading of the catalase over the various interfaces in monolayers or chemical inactivation by peroxide itself in reversible or irreversible reactions is responsible. These results may be compared with those obtained by George (65,66,67). I n this investigation a manometric technique was used which made it possible t o measure the rate of peroxide decomposition at constant peroxide concentration 0.01-5.0 M . There was an initial rapid evolution of oxygen which lasted for about two minutes, depending on the peroxide concentration, followed by evolution at a steady rate which slowly decreased in the course of an hour. This decrease was undoubtedly due t o enzyme destruction. I n the initial reaction the rate of oxygen evolution was found to decrease exponentially with time and this exponential decrease was more rapid the higher the peroxide concentration according to the following expression for the exponential constant:

where P is the peroxide concentration in moles per liter, and a, b, and c are constants having the values 0.072 M.-l sec.-', 0.15 M.-l, and 0.0185 M.-I sec.-l at O'C., respectively. Both initial and steady rates were directly proportional t o the catalase concentration but showed a complex variation with the peroxide concentration, the rates rising t o maximum values at 0.4 M and 0.07 M H202, respectively, and then decreasing markedly with further increases in the peroxide concentration. Peroxide itself thus inhibits its own decomposition in high concentration. These results were obtained with erythrocyte and liver catalase and lysed red cells, and further experiments showed that the po:isible presence of inhibitors in the peroxide or the nature and concentnition of the buffer solution played no part in the reaction. The inhibition of the steady rate decomposition at high peroxide concentrations was shown to be quantitatively reversible by several experiments in which dilute buffer solution or small amounts of concentrated peroxide solution were added during the course of the reaction. This inhibition is therefore not caused

T H E SPECIFIC REACTIONS O F I R O N IN SOME HEMOPROTEINS

395

by enzyme destruction. Dilution during the period of initial high activity also showed a reversibility which suggests that the transition from high to low activity is not caused by partial destruction of the enzyme. The initial high activity was found to be more susceptible to inhibition by azide and cyanide ions than the subsequent low activity for at certain concentrations of these inhibitors the initial high activity could be completely eliminated without affecting the low activity. Foulkes and Lemberg have confirmed that the activity of catalase decreases initially much faster than can be accounted for by its irreversible destruction but using different experimental conditions found that azide ions can enhance the initial falling off in activity (18). George (65) also investigated the kinetics of the decomposition of peroxide by concentrated catalase solutions in the presence of high concentrations of sodium azide (0.6 M ) and found a residual catalytic activity apparently attributable to azide catalase itself. The catalysis by azide catalase did not show the inhibition by peroxide at high peroxide concentrations like the free enzyme. The kinetics of the catalase peroxide reaction over a wide range of concentrations are thus very complicated and the problem is t o decide what are the significant kinetics for the oxygen evolution reaction by the intact enzyme. Bonnichsen, Chance, and Theorell (64) favor the view that only the first order kinetics given by Eq. (I) are significant. The bulk of their experimental data refer to peroxide concentrations of 0.1 M or less where George also found the initial rate directly proportional t o peroxide concentration. At higher peroxide concentrations, 0.3 and 1.0 M , data obtained by Millikan and McLaughlin (see 64) showed no decrease in rate at the higher concentrations in contrast to George's results. Further experimental work is very necessary in this concentration range to settle this point. The values for the velocity constant in Eq. (I) obtained by Bonnichsen, Chance, and Theorell for pure erythrocyte and liver catalase are k = 3.5 X lo7 and 3.0 X lo7 M.-l, set.-' at 22"C., respectively. The activation energy for the reaction with erythrocyte catalase has the remarkably low value of 1,700 & 100 cal. The comparison between the affinity of an inhibitor for catalase as measured spectrophotometrically and the affinity determined from kinetic measurements has revealed several interesting features. Keilin and Hartree (17) showed that azide ions and hydroxylamine are much more effective inhibitors than the affinity of these compounds for ferric catalase would suggest. Their results are given in Table VII, the first column giving the molarity of NaN3, NH20H, and KCN required to produce 50% inhibition, the second column giving the relative affinity

396

PHILIP GEORGE

TABLE VII Inhibition of Catalase Activity of Azide, Hydroxylamine, and Cyanide, the Relative A f i n i t y of These Substances for the Enzyme and the Corresponding Dissociation Constants for the Complexes Inhibitor

Molarity to Give 50% Inhibition

Relative Affinity for Ferric Catalase

IXssociation Constant of Ferric Complex

NaN3 NHiOH KCN

6.3 X lo-* 6.3 x 10-7 4.3 x 10-6

67 40 10,000

6 X 10-4 M. 10-8 M. 4 x 10-6 M.

of the three compounds for catalase determined spectroscopically, and the third column giving the dissociation constants of the azide ion and hydroxylamine complexes calculated from the dissociation constant of the cyanide ion complex as determined spectrophotometrically by Chance (68). If azide ions and hydroxylamine exercised their inhibitory power entirely by combination with ferric catalase, then the molarity required to give 50% inhibition would be numerically equal to the spectroscopically determined dissociation constant. The data shows them to be between lo3 and lo4 times more effective. Keilin and Hartree (17) have shown in a series of experiments that this enhanced inhibiting power of azide and hydroxylamine is connected with the reduction of catalase to a ferrous form when these substances are present, and they suggest that these compounds stabilize the ferrous form. The reduction of the iron is clearly demonstrated in the following way. When sodium azide is added t o a catalase c;olution, the color originally greenish brown becomes slightly more greenish and the absorption band at 622 mp is intensified and moves about 2 p nearer the blue end of the spectrum. Addition of hydrogen peroxide changes the color of the solution t o red, and the absorption bands are replaced by two stronger bands at 587 and 559 mp. If carbon monoxide is passed through the solution, these bands become more distinct and shift to shorter wavelengths, the band maxima now being at about 577 and 546 mp. Both these red compounds, especially the latter, are fairly stable in the absence of oxygen, but in its presence they rapidly change back into the original greenish brown azide catalase. Even in a nitrogen atmosphere enough oxygen is liberated from the hydrogen peroxide t o bring about this change although in a carbon monoxide atmosphere it occurs more slowly. Hydroxylamine catalase is also reduced by hydrogen peroxide to give a spectroscopically similar autoxidizable compound. A quantitative study of the inhibition of the peroxide decomposition by sodium azide in the presence of carbon monoxide showed that the

TEE SPECIFIC REACTIONS O F I R O N I N SOME HEMOPROTEINS

397

resultant inhibition was far more pronounced in CO/N2 than in C O / O n gas mixtures (Keilin and Hartree, 17). The data indicates that in CO/Oz mixtures there is a definite competition between the two gases for the ferrous azide catalase, the enzyme having a greater affinity for O2 than CO. Expressing the partition coefficient of azide catalase between 0 2 and CO in the form k =

[Cat. Oz]CO [Cat. CO1.02

where Cat. 0 2 and Cat. CO represent the concentrations of active and inactivated enzyme respectively, the experimental data were found to agree very closely with calculated values obtained using lc = 9. The inhibition of azide catalase by CO is very largely relieved by radiation from a mercury vapor lamp, which is additional support for the formation of a ferrous complex, for carbonmonoxyhemoglobin can be dissociated in this way, and it also confirms that very little irreversible inhibition is occurring. A further piece of evidence for the reduction of the iron in this system is the observation that although the affinity of cyanide for ferric catalase is about 150 times greater than that of azide (see Table VIIj, when cyanide is added to azide catalase, previously treated with hydrogen peroxide in a pure nitrogen atmosphere, no change in the spectrum occurs. The band at 587 mp remains sharper than the band at 559 mp in contrast to the cyanide-catalase spectrum in which the band at about 585 mp is broader and much less intense. Replacing the nitrogen atmosphere by carbon monoxide shifts the bands toward shorter wavelengths as was found in the absence of cyanide. Although these experiments establish beyond doubt the reduction of catalase hematin when azide or hydroxylamine are present, there is no similar direct evidence for reduction with the free enzyme, for in this system carbon monoxide has no inhibiting effect and no spectroscopic changes have been observed. The inhibition of catalase by cyanide shows none of the characteristics of the azide or hydroxylamine inhibition as is to be expected if cyanide combines with the ferric form. At low peroxide concentrations M the equilibrium constant for the formation of the cyanabout catalase complex (K1) determined from kinetic data using the expression

where R, is the uninhibited rate and R 1 the inhibited rate a t various (HCN) concentrations agrees well with the value of K I determined by direct spectrophotometric measurements (Chance, 68). This shows that under these conditions cyanide inhibits simply by the removal of

398

PHILIP GEORGE

active catalase as the inactive cyan-catalase complex :tnd that no competition occurs between cyanide and peroxide for the hematin iron atoms. This is in agreement with earlier observations that the extent of cyanide inhibition is independent of the peroxide concentration (see Lemberg and Legge, 5 ) . Simple though this appears, it is difficult t o correlate this observation with the formation of a catalase-hydrogen peroxide complex whose existence was revealed by a rapid-recording spectrophotometric technique developed by Chance (69). The results he obtained will now be described.

2. Complexes of Catalase with Hydrogen Peroxide and A1A:yE Hydroperoxides Catalase was found t o form an intermediate compound in the presence of hydrogen peroxide (Chance, 69). The spectrum was measured from 380-430 mp and is slightly shifted toward the visible as compared with free catalase. The complex shows no similarities to1 cyan-catalase or the compound formed when peroxide is added to azide catalase. Its formation is very rapid, the bimolecular velocity constant having a value of about 3 x lo7 M.-l see.-'. In the absence of added hydrogen donors, the complex decomposes slowly according to a first order reaction with a velocity constant of about 0.02 set.-'. This catalase complex thus resembles the green primary hydrogen peroxide complex of peroxidase. A very remarkable property of the complex is tha,t apparently only one of the four hematin groups of the catalase molecule is bound t o peroxide (Chance, 70). Three independent experimental methods provide qualitative evidence for this. First, the spectrophotometric titration of catalase with hydrogen peroxide gives amounts of the complex in excess of the molar amount of peroxide added when the calculation is based upon the combination of all the hematin groups with the peroxide. Secondly, the noncompetitive inhibition of catalase activity by cyanide indicates that not all the hematin groups are combined with peroxide under the conditions of the test for catalase activity. Thirdly, the difference between the extinction coefficient of catalase and the complex is rather small having a value of 40 cm.-l x mM.-l ak 405 mp which is comparable t o the difference between the extinction coefficients of erythrocyte catalase with four hematin groups (380 cm.-' X mM.-l) and liver catalase with three hematin groups (340 cm.-' X mM.-l). This small difference in extinction coefficient becomes even more significant when compared with the large difference found for the alkyl hydroperoxide complexes (180 cm.-' X mM.-' at 405 mp) whose formation and reaction will be described later. Although the peroxide decomposition by catalase shows noncom-

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

399

petitive inhibition by cyanide ions, Chance found that cyanide ions do compete with hydrogen peroxide in the formation of the peroxide complex (70). A quantitative study of the reaction of the catalase hydrogen peroxide complex with cyanide or alkyl hydroperoxides revealed that 1.2 +_ 0.1 hematin groups are combined no matter whether 1.3 of the 4 hematin groups are present as degraded bile pigments, and this composition is independent of large variations in the peroxide concentration. Chance suggests two possible interpretations of these results. If all the catalase hematins react in the same way and independently of each other, it follows that each group is saturated to the extent of (1.2 _+ 0.1) + 4 as a consequence of a steady state being set up, the complex being formed from catalase and peroxide and destroyed by reacting with further peroxide molecules. This interpretation, however, does not account for the constant composition of the complex when some of the hematins have been converted into bile pigment, as in liver catalase, without the additional postulate of an interaction between the catalase hematins. There is, however, no evidence for such interaction in the cases of combination with cyanide and methyl hydroperoxide which have been thoroughly examined for this effect (Chance 68,71). A second interpretation is that the peroxide molecule can only combine reversibly with one of the hematins of the catalase molecule and that having combined, the remaining hematin groups acquire “special properties ” such that they now react rapidly with peroxide and destroy it. This would account for a constant composition of one peroxide per catalase molecule independent of the bile pigment content, and thus conflicts with the observed value of 1.2 -t 0.1 bound hematins. More recent work by Chance and Herbert (72) on bacterial catalase has revealed that with this enzyme 1.6 peroxide molecules are bound to the catalase molecule. This suggests very strongly that the second interpretation entailing the unique combination with one hematin group is wrong and that the first interpretation in terms of a steady state is correct. Possible reaction mechanisms with the peroxide complex will be discussed later in comparison with hydroperoxide complexes. Like peroxidase, catalase forms both primary and secondary complexes with methyl and ethyl hydroperoxide (Chance, 73). The primary complexes are green having a diffuse absorption band in the red starting at 670 mp. The secondary complexes are red and have absorption maxima in the visible region at 572 and 536 mp. The catalase-ethyl hydroperoxide complex found by Stern (74) had maxima at approximately these wavelengths and was thus the secondary complex. The Soret bands of the primary complexes are similar in shape to that of the free enzyme but are shifted toward the red by several millimicrons. At

400

PHILIP GEORGE

405 mp both ethyl and methyl hydroperoxide complexes show a decrease of extinction coefficient as compared to the free enzyme of about 180 cm.-l X mM.-' for erythrocyte catalase, i.e., about 45 cm.-' X mM.-' per hematin group bound t o peroxide. The Soret bands of the secondary complexes are very different from those of free catalase and the primary complexes. They show a large shift toward the red, for instance the maximum for the methyl hydroperoxide complex is :st 422 mp with an extinction coefficient of 242 cm.-' x mM.-', and they appear somewhat similar to the cyan-catalase complex. A striking difference between the primary alkyl hydroperoxide complexes and the primary hydrogen peroxide complex is that in the former all the catalase hematins are combined (Chance, 70). I n the p H range 3.8-9.0 the combination follows the equation Cat. ( O H ) r

+ 4HOOR

Cat. (OOR),

+ 4EIoO

each group being bound independently. The bimolecular velocity constants for their formation although large are a little less than that for the formation of the hydrogen peroxide complex. The values for ethyl and methyl hydroperoxide and hydrogen peroxide are 2 X lo4, 8.5 X lo5, and 3 X lo7 M.-l set.-', respectively, showing a decrease in magnitude with increasing size of the peroxide. The formation of the secondary complexes from the primary alkyl hydroperoxide complexes is not simple (Chance, 71). The secondary complexes do not form until an appreciable amount of the primary complex is already present, yet the reaction is not a first order transformation of the primary complex as in the case of peroxidase. The velocity of formation of the secondary complexes increases with increasing alkyl hydroperoxide concentrations but not enough t o follow a second order equation. A property all the primary complexes have in common is the decomposition giving free catalase which does follow first order kinetics (Chance, 71). The velocity constants for ethyl and methyl hydroperoxides are 0.04 and 0.016 sec.-l as compared with 0.02 set.-' for the hydrogen peroxide complex. The secondary complexes decompose far more slowly, the first order velocity constants for ethyl and methyl hydroand 4 X SIX.-', respectively. peroxides having the values 2.3 X 3. Catalase in Coupled Oxidations

In 1936 Keilin and Hartree (75) showed that addition of ethyl alcohol t o certain enzymatic oxidation systems such as uricase with uric acid and amino acid oxidase with amino acids doubles the oxygen uptake.

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

401

Similar additions of ethyl alcohol to xanthine oxidase with hypoxanthine or xanthine oxidase with acetaldehyde had no effect; however, a little catalase added to these systems resulted in increased oxygen uptake like that observed in the previous cases. These results showed that alcohol undergoes a secondary or coupled oxidation by the hydrogen peroxide formed in the primary oxidation reaction. It was shown that catalase could utilize hydrogen peroxide for this type of coupled oxidation in a series of experiments in which hydrogen peroxide was gradually formed by the hydrolysis of barium and cerium peroxides. It therefore seemed likely that the difference between the primary oxidation systems was that uricase and amino acid oxidase contained slight traces of catalase whereas xanthine oxidase did not. Later experiments confirmed these results (Keilin and Hartree, 76), and the following set of compounds were found to undergo coupled oxidation: methanol, ethanol, n-propanol, isobutanol, 0-amino-ethanol, and ethylene glycol. Keilin and Hartree suggested that this coupled oxidation is a more likely biological function of the enzyme than the catalytic decomposition of the hydrogen peroxide into water and molecular oxygen. Chance (69,118) showed that this coupled oxidation can be accounted for by the primary hydrogen peroxide complex reacting with the alcohol in a bimolecular reaction liberating t,he free enzyme. The velocity constants for its reaction with methyl, ethyl, n-propyl, n-butyl, and isoamyl alcohols are lo3, lo3, 17, 2, and 0.1 M.-' sec.-l, respectively. Ascorbic acid was also found t o react with an apparent bimolecular velocity constant of 3.4 X lo2 M.-l set.-', but in a later paper Chance (119) has shown this to be a reaction of a different type involving the production of a secondary complex and not the free enzyme. I n addition in this paper, reactions of the primary complex with nitrite and formate are described which are similar to the reactions with alcohols. The primary compounds of the alkyl hydroperoxides also react with hydrogen donors. The methyl hydroperoxide complex reacts at approximately the same rate with methyl, ethyl, and n-propyl alcohol as does the hydrogen peroxide complex; the ethyl hydroperoxide complex reacts somewhat faster, the velocity constants for the three alcohols being 2.1 X lo3, 2.1 X lo3, and 33 M.-l see.-', respectively. The activity of the secondary complexes is negligible in dilute solution, so that formation of the secondary complex in the reacting system has the effect of inhibiting the reaction. This accounts too for the inhibition of the ordinary catalytic decomposition of hydrogen peroxide when ethyl hydroperoxide is added (Chance, 77). It is very interesting that hydrogen peroxide itself can react extremely rapidly with the primary alkyl hydroperoxide complexes in a manner analogous to the alcohols. The velocity constants for its reac-

402

PHILIP GEORGE

tion with the ethyl and methyl complexes are estimaied to be greater than lo6, and about 5 X 107 M.-l set.-', respectively. As in the case of catalytic decomposition of hydrogen peroxide the peroxidatic activity of the enzyme shows no inhibition by carbon monoxide. Chance investigated this in the system primary methyl hydroperoxide complex reacting with ethyl alcohol. Instead of inhibition a slight increase in the rate of disappearance of the complex was noted which could be attributed t o formate being produced by the hydration of the carbon monoxide and acting as an additional substrate (71).

4. Reaction Mechanisms with the Catalase-Peroxide Complexes Chance (78) has discussed this experimental data in terms of the extended Michaelis theory which accounts for the similar peroxidatic action of peroxidase, the only difference being that with peroxidase the main reaction can proceed via the secondary complexes, whereas with catalase these complexes are inactive and the main reaction proceeds via the primary complexes. Representing the primary complexes by FeOOH and FeOOR he suggests the various reactions are: (i) Oxidation of alcohol by the primary hydrogen per'oxide complex:

(ii) Oxidation of alcohol by the primary alkyl hydroperoxide complex:

+ -____.

FcO~I$- ROH +

E'c :67 OR ,

L

4

I n these two reactions the choice of the reactive oxygen atom in the peroxide group is purely arbitrary. (iii) Hydrogen peroxide reacting with the primary alkyl hydroperoxide complex : FerO)]OR ..-d

L--,

;H.!oo;Hj

+

FcOH -I- ROH 4

0:

+

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

403

(iv) Hydrogen peroxide reacting with the primary hydrogen peroxide complex, i.e., the “key” reaction in the catalytic decomposition of hydrogen peroxide : Fc :6: OH

.__a L--

f

IH.;oo!.H;

-

FeOH

-

E’COII + l i O B

+ €I:O + 0 2

Reaction (iii) as written is preferred t o the following alternative F c ; ? OR + 0

,

H i0.i OH

+

0 2

on the grounds that if this occurred an analogous reaction with an alkyl hydroperoxide would be possible : OR

FC

+ H j.Q OR 8

€I:O

,

+

-

FeOH

+ 2ROH + 0 2

and there is no evidence .at all for such a reaction evolving oxygen in the spontaneous decomposition of the catalase hydroperoxides. Important evidence in favor of reaction (iv) as the “key” reaction in the catalytic peroxide decomposition is the fact that the rate of formation of the primary hydrogen peroxide complex is high enough t o account for the entire catalytic activity, provided a mechanism for its rapid decomposition is available. The bimolecular formation velocity constant has a value of about 3 )( lo7 M.-l set.-' as compared with the value of the overall bimolecular constant for the catalytic reaction of 3 - 3.5 x lo7 M.-1 sec.-l. Any simple mechanism with consecutive reactions including the unimolecular decomposition must be excluded, for the decomposition rate of the primary complex is extremely low, the first order constant being about 0.02 set.-' (Chance, 69). The noncompetitive inhibition of the decomposition of hydrogen peroxide b y cyanide is not immediately obvious from the above reaction mechanism for if cyanide can compete in the formation of the peroxide complex which is responsible for the oxygen evolution in step IV, competitive inhibition might be expected. However, under the experimental conditions necessary t o observe peroxide decomposition, a n excess of peroxide is required and this is sufficient t o give the maximal concentration of the peroxide complex, 1.2 or 1.6 moles of bound peroxide for each erythrocyte or bacterial catalase molecule respectively, i.e., the peroxide complex concentration is independent of the peroxide concentration. Analysis of the system under these conditions shows noncompetitive inhibition t o hold.

404

P H I L I P GEORGE

Theorell (79) has suggested an alternative mechan-ism based on the combination of substrate and acceptor molecules with two hematin groups. Chance (78) has pointed out that this mechanism cannot apply t o the primary alkyl hydroperoxide complexes reacting with an alcohol or hydrogen peroxide (as in reactions ii an'd iii) because all hematin groups are attached t o hydroperoxide molecules in these complexes; however, it is applicable to the catalytic decomposition of hydrogen peroxide. The mechanism may be represented: HOFe-FeOH

HOFe-FeOOH

HOFe-FeOH

HOFe-FeOH

T

I

!

HOFe-IFe-OOH

-4 I HzOz

+ Hz0

HOFe-IFe-OOH

+HzO

+

0 2

The essential step evolving oxygen is the intramolecular reaction of the complex containing the two combined peroxide groups. There is no experimental evidence for this complex and for the mechanism to agree with Chance's data the complex would have t o have a very short half life. Chance has added that the following points are in favor of such a mechanism. First, it explains why peroxidase is a feeble catalyst for peroxide decomposition; with only one hematin group per molecule this type of mechanism is excluded. Secondly, it explains why methernoglobin is a poor catalyst, for methemoglobin only gives an inactive secondary peroxide complex. Thirdly, it accounts for the decrease in catalase activity with a decrease in the number of intact hematin groups. But this mechanism has the serious disadvantage that it cannot apply t o the reaction of the primary alkyl hydroperoxide complexes with hydrogen peroxide and in addition it requires further elaboration t o explain the peroxidatic action of catalase. Certain other aspects of these mechanisms will be discussed later in comparison and contrast with the reactions of ionic iron.

VI. THEMECHANISM O F THESEH E M O P R O T E I N REACTIONS A very valuable approach t o the problem of the general mechanism of these reactions has been made by Theorell (2) who has emphasized the manner in which structural differences between the various hemoproteins can be correlated with the appearance of highly specific reactions. An important feature of these particular structural groups is the interaction which exists between the reaction of the group and the specific reaction of the heme or hematin group. For instance, when oxygen is attached to hemoglobin in neutral solution, the pH is lowered

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

405

as a result of the ionization of a neighboring acid group. Theorell has named groups of this kind heme-linked groups and has examined the reactions of a wide range of hemoproteins for this type of interaction. There is not sufficient data yet to decide whether such groups play a primary role in the kind of oxidation-reduction reactions discussed above and undergo oxidation-reduction themselves, or whether their function is secondary in that by ionization the groups affect the speed with which the heme or hematin group reacts. I n the case of the hemoglobin, a mechanism will be advanced later in which the assumption of a n auxiliary electron accepting group on the protein molecule makes it possible to explain the autoxidation kinetics. But first it is necessary t o examine the several mechanisms which have been suggested to explain the various oxidation-reduction reactions. It has been shown above that the catalytic action of catalase and peroxidase is intimately connected with the ability of these hemoproteins to form complexes with hydrogen peroxide (or alkyl peroxides). By choosing the experimental conditions the existence of three different complexes can be demonstrated spectroscopically. The chemical nature of these complexes is as yet unknown, and mechanisms have been represented as bimolecular reactions between substrate and complex. It is difficult to reconcile such mechanisms in which some four strong covalent bonds are broken and then re-formed simultaneously with the high speed and low activation energy of these reactions (Bonnichsen, Chance, and Theorell, 64), and some stepwise mechanism seems more likely. These kinetic features in the case of oxidation-reduction systems are usually associated with very simple reactions such as electron transfer or reactions in which only one bond is broken and another formed. The reactions of hydrocarbon free radicals, as elucidated by the “sodium flame” technique developed by Polanyi and his school (80) and the part they play in the oxidation of hydrocarbons as evidenced by many kinetic studies are cases in point (Pease, 81; Steacie, 82). The mechanisms proposed by Lemberg and Legge ( 5 ) have a n added disadvantage. An acceptor group joined to the protein or the hydrogen accepting substrate linked to the protein is made to react with HzOzor OzH- linked to the hemoprotein iron atom by an intramolecular rearrangement. With the reacting groups attached in this way some distance apart, it is difficult t o see how a reaction involving the breaking of several bonds and transfer of atoms or fragments of molecules can occur, for the gain in energy coming from the partial formation of the new bonds, which can normally offset a large activation energy needed to break the old bonds, would appear to be excluded. Some of these difficulties may be illusory if the structure of the or [Fe,-H202]+ but is peroxide complex involved is not Fe,-OOH

406

PHILIP GEORGE

somewhat simpler (Fe,+ represents the ferriprotoporp hyrin group joined t o the protein). The existence of three complexes with very distinct spectroscopic characteristics formed from the same two molecules suggests that one or two of these are reaction products or degradation products of the two components. At the most two might be accounted for by Fe,-OOH and [Fe,-H20z]+. There is no evidence that any of these complexes has involved an irreversible reaction with a group on the protein, and it is not easy t o see how any such reaction should alter the absorption spectrum so markedly in the visible region. Likewise there is no evidence for irreversible attack on the porphyrin ring itself. The experiments of Chance show the interrelation of two of the complexes. The green complex of peroxidase (complex I) changes t o the red complex (complex 11) according t o first order kinetics. It seems highly likely then that the possibility of a complex itself undergoing a reaction must be taken into account in a complete reaction mechanism, although under some experimental conditions such a reaction may not play a dominant role. In its wider imp1ic;ations this means that the entire question as to whether a valency change occurs and the conjugate problem of the participation of free radicals as intermediates must be reconsidered. I n discussing this a comparison has often been made between hemoprotein reactions and those of ionic iron (Stern, 3). Calculations have been made showing that the relative activity of catalase, hemin and ionic iron as catalysts for hydrogen peroxide decomposition are in the ratio lo6: 10P. However, such calculations can be misleading, for the reactions show different dependencies on the hydrogen ion concentration which make a valid comparison impossible. The 10l0-fold greater efficiency of catalase as compared with ionic iron, which these figures show, could be taken t o suggest an entirely different type of reaction mechanism. Yet if ionic iron were not precipitated in alkaline solution, a t pH 10.46 it would have a catalytic activity equal to the maximum activity of catalase itself. This conclusion follows from the values for the velocity constant k in the rate equation. -dHzOz/dt

=

k[H~O~][Catalyst]

which for catalase and ionic iron (neglecting the formation of hydroxides) are 3.5 x 107 and I .2 x 10-3/[H+]M.-1 set.-', respectively. For catalase and peroxidase there is no direct evidence that a valency change of the iron atom occurs except when azide js added to catalase, and peroxidase is acting as an oxidase in the autoxidat ion of dihydroxymaleic acid, when both systems can be poisoned by carbon monoxide. The normal absence of such poisoning cannot be taken as proof that the

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

407

iron remains in the trivalent state throughout the reaction, for an alternative explanation is that although a valency change occurs, other reactions competing for the ferrous form are more rapid than the combination with carbon monoxide. In this connection the recent experiments of Keilin and Hartree (45) with methemoglobin and hydrogen peroxide are significant, for the system shows a well-developed peroxidatic and catalatic action and reduction of the iron is apparent from the formation of HbOz, orHbCO when carbon monoxide is admitted. In a like manner the different oxidation products formed when peroxidase and ferrous ion are used in peroxidatic reactions (109,110,111) do not prove that the underlying mechanism is different. The mechanism in the case of ferrous ion involves competition between Fe++ and Fe+++ for radical intermediates, and it is quite possible with reactions of such complexity, particularly if they proceed through several distinct oxidation stages, that such competition reactions could lead to different reaction products predominating, depending on the catalyst employed. In the absence of direct proof of a valency change or free radical intermediates participating, it is valuable to survey the data on the reactions of ionic iron in similar systems and then examine the possibility of similar mechanisms occurring in hemoprotein reactions.

1. Free Radical Mechanisms with Ferrous and Ferric Ions The evidence for the free radical mechanisms of the reaction between ferrous and ferric ions and hydrogen peroxide is fully discussed in the article by J. H. Baxendale in this volume, and it is necessary here only to summarize and comment on those features especially relevant to hemoprotein reactions. This evidence is essentially indirect. Experiment shows very reactive intermediates t o be present and extensive kinetic studies reveal competition reactions for these intermediates in that the overall order of the reaction is found t o depend on the reactant concentrations. A free radical mechanism is adopted because it accounts for the chemical reactivity of the system in the oxidation of substrates (Fenton's reaction) and the initiation of the polymerization of vinyl compounds (Baxendale, Evans, and Park, 84) and it provides a set of reactions which largely account for the observed kinetics. The set of reactions which fit best the most recent experimental data is that proposed by Barb, Baxendale, George, and Hargrave (83) : Fe3+ Fez+ Fez+ HO Fez+ Fe*+

+ OzH- Fea++ HOz + Hz02 Fe3++ OH- + HO + HO Fe3++ OH+ HzOa H a 0 + HOz + HOn + Fe3++ OZH+ + Fez+ + O2 + -+ 4

+

02-

(i') (0) (1) (2)

(3)

(4')

408

PHILIP GEORGE

(Reactions i and 4 would refer to corresponding reactions with un-ionized

HzOzand HOz.)

This reaction scheme is to be preferred to the original mechanism proposed by Haber and Weiss (85) in which oxygen was evolved in the step HOa

+ Ha02 -+

0s

+ Ha0 + HO

(w)

for two main reasons. The Haber and Weiss mechanism is incompatible with the recent kinetic data, and it does not provide a consistent mechanism for the reaction of ferrous salts with hydrogen peroxide and the catalytic decomposition of hydrogen peroxide by ferric salts. George’s stoichiometric experiments with potassium superoxide, KOz, also indicated that reaction w is insignificant when ferric and ferrous ions are present in solution (86). Although superficially similar in so far as some reactions occur in both mechanisms the kinetic differences between the two are fundamental. In the Haber and Weiss mechanism the ferrous and ferric ions only initiate and terminate a reaction chain propagated by steps w and 2 above which involve the two chain carriers, the HO and HOz radicals. I n the new mechanism there are three chain carriers, the radicals HO and HOZ and the ferrous ion itself. The new mechanism is supported by the elucidation of the thermochemistry of the various steps in terms of the ionization potential ( I ) of the ferrous ion in aqueous solution, the electron affinities ( E ) of the radicals HO and HOz plus the heats of solvation (8)of the corresponding ions OH- and OSH- in water, the various 0 * H arid 0 * * 0 bond strengths ( D ) and other thermal quantities such as heats of evaporation (A) and heats of solvation (8) (see 83,84,87,88,89). A411the steps with the exception of (i’) and (0) are exothermic, and being; electron transfer reactions or a simple bond breaking reaction (step 2) are to be expected to proceed extremely rapidly as is required by the kinetic mechanism. Steps (i’) and (0) are endothermic to the extent of 28 and 5 kcal., respectively, and these values are in accord with their activation energies which are 28 5 8 and 9.4 kcal. As with the kinetic evidence this thermochemical evidence is indirect support for the radical mechanism. An important point which will be elaborated later is that thermal data of this kind provide criteria for rejecting certain steps when the endothermicity becomes impossibly large for the reaction to proceed a t a sensible rate. There are some further aspects of the ionic iron-hydrogen peroxide system which have a possible bearing on hemoprotein reactions. Apart from the reactions of the HO radical listed above (1 and 2) and reactions with oxidizable or polymerizable substrates, there is Eome experimental evidence that two additional types of reaction may occur. The first

- -

-

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

409

is electron transfer from anions, e.g., HO.

+ C1- --+

OH-

+ C1.

a reaction which is energetically favored. The second, and perhaps the more significant, is a reaction with the ferric ion itself. The evidence for this is not so clear cut as in the previous case, yet it is difficult to see any other interpretation for the observation that in the presence of higher concentrations of ferric salts than those usually employed (10-1 M rather than M ) the ratio of the velocity constants of steps 1 and 2 has a different value, control experiments having shown that a normal kinetic salt effect was not responsible. Except that it takes part in reactions similar to those of the HO radical, there is no evidence as to the nature of this ferric ion-HO radical reaction. One possibility is the formation of a ferryl ion FeO++, a compound of quadrivalent iron: FeS+

+ HO

--t

FeOHS+ -+ Fe02+

+ H+

The participation of the ferryl ion in the ionic iron-peroxide system was suggested in 1932 by Bray and Gorin (90) although its relationship to, and possible role in replacing, the HO radical was not realized. Cupric ions which for some time have been known to accelerate the ferric ion catalyzed decomposition of peroxide (Bohnson and Robertson, 91) have also been shown to affect the ferrous ion reaction (83). There is now good kinetic evidence that this arises from the cupric ion reacting more rapidly than the ferric ion with the 02-: cu2+

+ oz--t

Cu+

+ Fe3+-very fast

followed by

cu+

+

01

Cu2+

+ Fez+

The second step accounts for the fact that the action is purely catalytic and that a limit can be reached above which further additions of cupric ion have no effect (Barb, Baxendale, George and Hargrave, 83). One point of marked resemblance between ferric ions and the hemoproteins in their ferric state is the formation of a complex with hydrogen peroxide. By the appropriate choice of peroxide and hydrogen ion concentrations a deep brown-red colored complex is formed with ferric ions. Evans, George, and Uri (88) showed this to be an “ion pair” complex formed: FeS+

+ OzH-

[FeOZHla+

The heat and entropy of formation of this complex have values in accord with the data on other ferric ion pair compiexes, e.g., FeOH2+ and FeCI2+. In contrast with the apparent behavior of the hemoprotein-

410

PHILIP GEORGE

hydrogen peroxide complexes this ferric ion complex plays no distinct part in the reaction mechanism. A bimolecular reaction between the complex and another hydrogen peroxide molecule does not occur. The kinetics of the peroxide decomposition can be explained if the initiating reaction is (i') in the above set of reactions, and there is no means of telling from the kinetics whether reaction (i') occurs by unimolecular decomposition of the ion-pair complex or simple collisions between Fe3+ and OzH-, or both (83). Recent studies have shown that the oxidation of a substrate by the ferrous ion-hydrogen peroxide system is a reaction of great complexity. I n the absence of oxygen the following general scheme offers an explanation for the chief kinetic features. Kolthoff and Medrtlia (92,93) found it to hold for the oxidation of ethyl alcohol to acetaldehyde and Barb, Baxendale, George, and Hargrave (83) found it to apply to the oxidation of traces of organic impurities in distilled water. H.ZA represents the oxidizable substrate.

+ + + + + + +

+ + + + + + + +

Fez+ HZOZ + Fe3+ OHHO Fez+ HO + Fe3+ OHHO+ HA HzO HzA Fe3++ HA+ Fez+ HA HA+ OH-+ HAOH HAOH HO + HOA HzO Fe3+ A 0 Fez+ H+ HOA --f

The presence of oxygen profoundly affects the reaction. This is similar to the effect of oxygen on the initiation of the polymerization of vinyl compounds by the ferrous ion-hydrogen peroxide system (Baxendale, Evans, and Park, 84). It is due t o the addition. of oxygen t o the organic radical, and when oxygen is present in these oxidation systems the following reactions have also to be taken into account (84,92,93).

+ -+

HA 0 2 + HAOz HAOz Fe2+ H+ -+ HAOzH Fea+ Fez++ HA0 4-Fe3+ OHHAO,H HzA+ HAOH HA HA0 HA02 HzA + HAOzH HA HA0 Fez+ H + + HAOH Fe3+

+

+

+ + + +

+ + + + +

The use of ferric ion and hydrogen peroxide as an oxidizing system has been very little studied. This is surprising because it provides the simplest model for peroxidase. Walton and Christiansen (94) showed that ethyl alcohol could be oxidized at 30°C. and the reaction was retarded in more acid solutions. The apparently low catalytic activity of this system would appear to be due t o the slowness of the initiating

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

411

reaction (i) arising from its big activation energy Fe3+

+ H2OZ

Fe2+

4

+ HO, + H+

(i)

compared with the corresponding reaction with the ferrous ion which is rapid Fez+

+ H2OZ+ Fe3+ + OH- + HO

(0)

At 25°C. and pH 2.0 the velocity constants ki and k, have the values 9.1 X and 53.0 M.-l sec.-l, respectively. For the reactions with hydrogen peroxide there thus exists a large accumulation of experimental data t o compare with the data on hemoprotein reactions. However, this is not the case for the reaction with molecular oxygen where the experimental data is very scanty. Weiss (95) has suggested the following radical mechanism for the autoxidation of ferrous salts and this has been generally accepted.

+ O2 Fe3++ 02+ H+ HOz 3 Fez+ + HOz Fe3+ + OzHOzH- + H+ 2 HzOz 2Fe2+ + HZO2 2Fe3+ + 20Ha

Fez+

4

02-

-+

-+

From the appropriate stationary state equations it can be shown that the oxidation rate should be:

and initially when Fez+>> Fe3+ -dFeZ+/dt = 4k,FeZ+02

The experimental data of Lamb and Elder (96), however, are not in agreement with this predicted rate expression for they find the initial rate proportional t o the square of the ferrous ion concentration and directly proportional t o the oxygen pressure. This has recently been confirmed by the author (97), and it would appear that either the autoxidation is subject t o a true catalysis by trace impurities (an induced reaction is excluded by the total ferrous ion oxidized being large, about M / 2 0 ) or the actual mechanism is different from t h a t suggested by Weiss. 2. An Examination of the Possibility of Similar Free Radical Reactions

with Hemoproteins The purpose of this discussion is t o see which of the reactions listed above for the free ferrous and ferric ions are possible with the hemo-

412

PHILIP GEORGE

proteins. An examination of some of the experimental data on these reactions to try to decide whether they do in fact occur will follow in the next section. In considering the thermochemistry of the above reactions the only quantity which changes when ionic iron is replaced by another metal or coordination complex is the ionization potentisil I of the reduced form. The following list gives the heats of the reactions (i’--4’) in terms of electron affinities, solvation energies, bond strengths, etc., as given above, using the most recent value for these quantities (83,87,88,89).

+ OzH-+

+

Fez+ HOZ +I - ( E H o ~ SO~A-)S H O=~+I - 123 kcal. Reaction (0) : Fez+ HzOz -+ Fe3+ OHOH Qo = - I - X H ~ O , - S H ~-ODEO-OH ~ (EHO SOH-) SHO== - I Reaction (1): Fez+ OH + Fe3+ OH-

Reaction (i’): Fe3+ Q,r

=

+ +

+

+

+

+

+ +

+

+

+ 89.7 kcal.

- I + (EHO + SOH-)- SHO= - I + 137.6 kcal. + HOZ + Fe3+ + OZH= - I + ( E H o+ ~ SOZH-) - S m a = - I + 123 kcal. + Fez+ + Reaction (4’) : Fe3+ + Qr I - (Eoz + Soz-) + Soz = I - 76 kcal. Fe3+ + Reaction (a): Fez+ + QB = - I + (Eoz + Soz-) - So2 = - I + 76 kcal. QI

=

Reaction (3): Fez+ Q3

02-

0 2

=

0 2 -

02-

Unfortunately there is not sufficient thermal data for direct calculation of I in the case of heme, hemochromogens, and reduced peroxidase. In the first three cases, however, estimates can be based on a knowledge of their oxidation-reduction potentials. The ionization potential I for a metal complex M2+ionizing in solution is given by the heat of the reaction: and since

I can be obtained from the heat of the cell reaction ( - A H 0 ) M(aq.)’+

+ H++

M(aq.)’+

+ 3Hz - A H o

from the equation: - A H o = 82 - I

or I = A H o

+ 82 kcal.

Knowing the standard free energy change in the cell reaction, A H o can be calculated if ASo is known. From standard entropy values the change Hf to &Hz corresponds to f15.6 e.u. and so only a n estimate of the entropy change associated with the valency change of the complex is required to give a value for A H o and hence I .

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

413

In'the case of such large ions as the metalloporphyrin complexes with their rigid planar structure, t o a first approximation the entropy change will be given by the difference in the Born charging entropy of the two valency forms. The Born charging entropy of an ion is given by: 8B.C. =

-2R D 2

where Z is the magnitude of the ionic charge, e the electronic charge, R the radius of the ion, D the dielectric constant of the medium, and ( d D / d T ) , the temperature coefficient of the dielectric constant at constant pressure (Eley and Evans, 98). Taking heme as a sphere of radius 5.8 A. equivalent in volume to its platelet size of 15 X 15 X 3.7 A.3, the change in entropy from the reduced to oxidized form amounts to - 1.7 e.u., considering only the central iron atom (i.e., change in charge 0- +1) and +5.0 e.u., taking into account the ionized propionic acid side chains of the protoporphyrin (i.e., change in charge - 2 + -1). The total entropy change in the cell reaction is thus 13.9 or 20.6 e.u., which give for TAXo, 4.1 or 6.1 kcal. As the difference between these values is small considering the assumptions made in the calculation, a n approximate value of TASO of 5.0 kcal. has been used in the calculation of I listed in Table VIII, which also includes values of I for Fe2+, Co2+,Cr2+, Cu2+and Fe (CN)64-for comparison. TABLE VIII Calculations of Ionization Potentials in Aqueous Solution Complex or Metal Ion Heme.. . . . . . . . . . . . . . . . .. a Cyanide hemochromogen . . a Nicotine hemochromogen . .Q Globin hemochromogen . . .a Hemoglobin. . . . . . . . . . . . . .b Myoglobin . . . . . . . . . . . . . . .c Fez+. . . . . . . . . . . . . . . . . .d coz+. . . . . . . . . . . . . . . . . . e Cr2+. . . . . . . . . . . . . . . . .. * c u * + .. . . . . . . . . . . . . . . . .I Fe(CiY)sa-. . . . . . . . . . . . . I

Eo OT Eo'

AG

Volts

ICcal.

+0.110 +O .183 -0.184 $0.098 -0.144 -0,046 -0.77 -1.84 + O . 41 -0.167 -0.47

- 2.5 - 4.2 4.2 - 2.3 3.3 1.1 +17.8 $42.5 - 9.5 3.8 +10.8

+ + +

+

TASO Kcal.

Kcal.

+ 5.0 + 5.0 + 5.0 + 5.0 + 5.0 + 5.0

+ 2.5 + 0.8 + 9.2 + 2.7 + 8.3 + 6.1

-

4.9 4.9 4.9 7.4 +14.2

A€€O

$12.9 +37.6 -14.4 - 3.6 +25.0

I Kcal. 84.5 82.8 91.2 84.7 90.3 88.1 94.9 119.6 67.6 78.4 107.0

~

a

ED'values at pH 7.0 and 30'C. (Guznian Barron, 99).

Taylor and Hastinga (100). Taylor (101), Wyrnan and Ingalls (102). Taylor and Morgan (103). Evans, Baxendale, and Uri (89). * Eo values from Latimer (104). Entropy changes assumed to be identical with Fez+ systems. I EDand entropy of Cu*+from Latimer (104). Entropy of Cu+ assumed identical with Na+. 0 Data given in Butler (105) corrected for AH0 = -82 kcal. for H + e+ 1Hz (89). In a, b , c the free energy changes are not strictly standard free energy changes, for the redox potentials refer to unit concentrations and not unit activities of the reactants. b

+

414

PHILIP GEORGE

The only available check with experiment for the heme compounds is in the case of hemoglobin where a t ph. 7.0, Eo' = -0.152 at 24°C. (Conant and Pappenheimer, 106) and -0.144 or -0.14!3 (Taylor, 101) at 30°C. A H o of the cell reaction is thus 12.6 or 6.9 kcal. and the corresponding value of I , 94.6 or 88.9 kcal. The calculated value is in tolerable agreement with these experimental values, but it is clear that a thorough investigation of the variation of Eo' over a much wider temperature range on the same sample of hemoglobin and using the same experimental method is needed to establish a reliable value. The fairly close agreement between calculated and experimental values does show, however, that the assumption regarding the entropy change awociated with the valency change is not far wrong, and the entropy change is in fact small. It would appear from this that no marked reorganization of the protein structure accompanies the valency change. These calculations of the ionization potential show that heme, hemochromogens, and hemoglobin should in their reduced form react as rapidly or more rapidly than ferrous ions. Reactions (3) and (4) are again exothermic and should be very rapid. Reaction (0)is about thermoneutral instead of 5 kcal. endothermic, and so this direct reaction of the ferrous form with hydrogen peroxide should be faster than with ferrous ion itself. Reaction (a), a step to be considered in an autoxidation mechanism, is also less endothermic than in the case of ferrous ion and again a higher rate is to be expected. Reaction (i'), between the ferric form and OzH- is conversely more endothermic and should be much slower than with ferric ions because of this and also because the temperature independent factor should also be less than lo2*,the value found for the ferric ion reaction. This high value is attributable to reaction taking place between oppositely charged ions, one carrying a high charge. If this reaction is considered as taking place between the ferric form and the H202 molecule (instead of the OzH- anion) it will be 8.2 kcal. more endothermic on account of the heat of ionization of the hydrogen peroxide. A value of the bimolecular velocity constant for a hemin derivative reacting with hydrogen peroxide at ph. 7.0 and 20"C., assuming I = 90 kcal., the activation energy equal to the endothermicity and a normal temperature independent factor of lo", is given by: k =

1O1I . exp . (90 - 123) X lo3/&" X 1.8 X lo-'* = 3.6 X 10-7

where 1.8 X 10-l2 is the dissociation constant of hydrogen peroxide. I n spite of the very approximate nature of this calculation arising from uncertainty as to a good value for the temperature independent factor there can be no doubt that the magnitude of the velocity constant will

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

415

be very small indeed, and it is very unlikely that this type of reaction can play any significant part in heme or hemoprotein reactions. This is especially true in the case of catalase which has not been mentioned in this section because all attempts to prepare the ferrous form have failed. The most likely reason for this is that its ionization potential is very low and the ferrous form is able to react with any oxidizing entity present, even H+ ions. It may be concluded from this discussion that heme, hemochromogens, and the hemoproteins could undergo similar free radical and electron transfer reaction as the free ferrous and ferric ions. On thermochemical grounds reactions (l), (3), and (4)should be very rapid and reactions (0) and (a) should be more rapid than with ferrous ion in contrast to reaction (i), the ferric form reacting with hydrogen peroxide which should proceed extremely slowly. 3. The Applicability of These Free Radical Mechanisms in the Hemoprotein

Reactions a. Reactions of Peroxidase and Catalase. The observation (Dainton and Smith, 107) that neither catalase nor peroxidase can initiate the polymerization of vinyl compounds in the presence of hydrogen peroxide cannot be taken as evidence that free radicals are not formed when they react with peroxide. This observation resembles the absence of carbon monoxide poisoning in the reactions of peroxidase and catalase. It could be that other reaction steps compete for the radicals in the first case and the reduced form of the enzyme in the second case to the exclusion of the test reactions with vinyl compounds or carbon monoxide and so lead to a negative result. However, the above estimate of the heat of the reaction when a hemin derivative reacts with hydrogen peroxide in a manner analogous to the ionic ferric ion reaction: Fe3+

+ OzH-

+

Fez+

+ H01

makes it extremely unlikely that a reaction of this type is the initial redox reaction with peroxidase and catalase. The primary and secondary complexes that these enzymes give with peroxide appear to play the part of intermediates in an alternative reaction sequence by which the peroxide molecule undergoes reduction. In this respect this process is diametrically opposed to the ionic ferric ion reaction above which is the first step by which the peroxide molecule can undergo oxidation to molecular oxygen. Chance has not discussed in detail possible structures for the primary and secondary complexes except to point out that the Soret bands of the primary complexes with catalase resemble those of open chain porphyrin degradation products like degraded hemoglobin and

416

PHILIP GEORGE

specimens of catalase containing a high bile pigment content. There can be no doubt that peroxide combines directly with the iron atom of the hematin group and he suggests that “speculation a,s t o whether the porphyrin ring is actually oxidized on formation of the primary complex by electron transfer from the iron peroxide complex and is then reduced on reaction with the reducing substrate or acceptor, affords very interesting possibilities” (73). I n the absence of direct evidence all arguments a t present as t o the nature of these peroxide complexes can only be based on analogy with other hemoprotein derivatives and their reactions. Chance’s suggestion t h at the primary catalase complexes involve attack on the porphyrin ring does not seem very likely when comparison is made with peroxidase. The primary peroxidase complex resembles the primary catalase complex particularly in the Soret region, but the structure with a modified porphyrin ring is attributed to complex I V formed with peroxidase in the presence of excess alkyl hydroperoxide. I n the visible region the spectra of the primary complexes resemble the free enzymes, the very rapid formation of the complexes, and in the case of catalase the competitive formation in the presence of cyanide, all suggest that the formation involves a simple bimolecular displacement. If this is so then either H202or 02H- is joined to the iron atom. The high affinity of the hematin group of the enzymes for hydrogen peroxide tends t o exclude the bonding of the H202 molecule as such, for this type of reaction is in the category of replacement of a bound OH- anion by H202 of solvation and this reaction would appear unfavorable both on considerations of heat and entropy changes. Bonding of the OzH- anion is more likely t o account for the high affinity particularly if an OH- anion and not a H 2 0 molecule is to be replaced. It is interesting to note t h a t a compound of this type is not restricted to hemoproteins. The complex obtrained with ionic ferric iron is of this type having an “ion pair’’ structure Fe02H2+(Evans, George, and Uri, 88). Nevertheless there is an important distinction in the function of such complexes in the two systems. With ionic iron the compl ex plays no independent kinetic role, electron transfer predominates and there is no indication whether it occurs by unimolecular fission of t h e complex or direct bimolecular reaction hetween Fe3+ and 02H-, but the net result is one step toward the oxidation of the peroxide molecule. With the hemoproteins the peroxide complexes play a dominant kinetic role connected with a direct overall reduction of the peroxide molecule. It is true that with excess peroxide catalytic decomposition occurs in all cases which may be regarded as a mutual oxidation reduction of one molecule by the other. The distinction in the two cages is in the fate of the peroxide molecule involved in the complex.

THE SPECIFIC REACTIONS

OF IRON IN SOME HEMOPROTEINS

417

The conversion of the green primary complex into the pale red secondary complex appears t o be a reduction process even though it occurs in the absence of any added hydrogen donors. The most definite evidence for this is the case of peroxidase where the speed of the conversion is increased in the presence of all compounds with which the peroxide system reacts (Chance, 5 5 ) . For catalase, where the conversion can only be obtained with alkyl hydroperoxides, the evidence is not so clearcut, but a t least the velocity of formation of the secondary complexes increases as the hydroperoxide concentration is increased. An alternative explanation for these effects would be that the primary and secondary complexes are in some sort of equilibrium where removal of the latter would have the effect of increasing the rate of conversion. There is no indication of any such equilibrium, however, and direct reduction of the primary complex appears t o be the most likely explanation. One possible formulation for this change involves the production of a “ferry1 ion’’ type of compound by the removal of an OH radical by the hydrogen donor from the OzH- anion bound t o the iron atom: Prot Fe,OOH

+ AH2-

Prot Fe,O

+ HzO + AH.

(green primary complex) (pale red secondary complex)

The “ ferrylperoxidase ” being nominally a compound of quadrivalent iron would have oxidizing properties corresponding to one equivalent. This raises a difficulty encountered in any attempt to formulate a more detailed reaction mechanism taking account of the increased rate of formation of the secondary complex in the presence of the hydrogen donor. Chance (54)has shown that the secondary complex when formed takes part in a bimolecular reaction involving two equivalents of the hydrogen donor. As hydrogen peroxide initially only has an oxidizing capacity of two equivalents, this is incompatible with the conversion itself being a reduction process. * Further experimental studies are required to resolve this difficulty. In view of the complications which can occur with the ferrous ion-hydrogen peroxide system when it is used as an oxidizing agent in the presence of molecular oxygen, a test of the effect of removing oxygen from these enzyme reactions appears desirable. It is abundantly clear that further discussion of the possible free radical nature of subsequent reactions in these enzyme systems must await a more detailed knowledge of the structure of these complexes. The recent experiments by Keilin and Hartree (45) showing most clearly that methemoglobin and metmyoglobin are reduced when they

* This difficulty may only be apparent, for the bimolecular reaction was largely studied in the presence of an excess of hydrogen donor where the kinetics become pseudofirst order and the underlying stoichiometry can be obscured.

418

PHILIP GEORGE

react with hydrogen peroxide in sufficiently high concentration to cause catalytic decomposition, taken in conjunction with their previous demonstration (1G,17) of the reduction of azide catalase by hydrogen peroxide, are indisputable evidence that a reaction path for peroxide decomposition by hemoproteins involving a change of valency does exist. It is difficult, however, to formulate a mechanism for the methemoglobin and metmyoglobin reactions because again a peroxide complex of uncertain structure takes part. The experiments of George and Irvine (1 16) on the metmyoglobin complex showed t h a t it is probably formed in a reaction of the type MetMb

+ H 2 O 2 + Complex + X

(2 oxid. equiv.) (1 oxid. equiv.) (1 oxid. equiv.)

and that a transient oxidizing entity X is produced along with the complex, which can react with added reducing agents if present initially, but which disappears if the complex formation is allowed to proceed independently. A tentative mechanism which can account for these results is the formation of an OH radical and a “ferTy1’’ compound of myoglobin in the reaction i.e.,

+ Prot Fe,O + OH + €LO+ + H 2 0 2 + Complex + X

Prot FeP+(H20) MetMb

H202+

The reduction of the iron t o the ferrous form in the presence of excess hydrogen peroxide which was observed b y Keilin and Hartree (4.5) could then occur by ferrylmyoglobin reacting in a simple bond breaking reaction Prot Fe,O

+ 02H-

Ferrylmyoglobin

-+

Prot FepO*

+ OH-

Oxymyoglobin

The same reaction in the case of peroxidase would suggest that the deep red complex peroxidase-H202 (111) has the oxyheinoglobin type of structure. The participation of ferry1 compounds in this manner offers an alternative way for a ferric compound to react in a series of simple steps comparable t o the simple free radical reactions. In a more recent series of experiments, George (97) has shown that the secondary hydrogen peroxide complexes of horseradish peroxidase and cytochrome-c pcroxidase can be titrated with ferrocyanide or ferrous ions and alvo appear to take part in a one oxidizing equivalent reduction t o the ferric form of the enzyme. In this important chemical property they thus rcsemblc the metmyoglobin complex in spite of marked spcctroscopic differences in the visible region of the spectrum (Keilin and Hartree, 48). If the peroxide molecule is not a component part of the structure i t

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

419

was to be expected that other oxidizing agents might, be found to yield the secondary complex; HOCl, HOBr, Br03-, IOa- and (310, form a compound spectroscopically indistinguishable from the secondary H z 0 2or MeOOH complexes of horseradish peroxidase. Cytochrome-c peroxidase forms its secondary complex with all but HOCl and HOBr, which cause degradation of the enzyme. In a search for other redox systems capable of reacting with either the free enzyme or the secondary complex it was found that chloriridite ions reduce both secondary complexes whereas ferrous tris-dipyridyl ions or ferrous tris-ortho phenanthroline ions do not. This indicates that the redox potential for the reaction Complex I1

+ electron

4

Ferric form of the enzyme

lies between -1.02 and -1.06 volts, these being the potentials for the chloriridite ion and ferrous tris-dipyridyl ion couples respectively. These experiments suggest that the secondary complexes should no longer be regarded as Michaelis-Menten enzyme substrate complexes but as reaction intermediates in the same sense that free radicals and semiquinones are reaction intermediates, for all three classes of compounds provide a path for stepwise reactions. As a consequence the accepted mechanism for peroxidase action needs revision. There are several structures which may be considered for the metmyoglobin complex and the secondary complexes of the peroxidases: a) A simple quadrivalent ion structure containing Fe4+ b) A derivative of quadrivalent iron such as the ferry1 ion, Fe02+(see George and Irvine, 116, 120) c) A diradical structure in which trivalent iron forms “one end of the radical,” the other end being a normal radical grouping at a methene carbon atom, a pyrollic carbon atom, a some other atom within the conjugated network of porphyrin ring and heme-protein linkage d) A higher oxidation state of the hematin group in which the electron has been removed from a n-orbital common to the ring as a whole. The existence of this type of oxidation state has recently been demonstrated by Cahill and Taube in the case of the structurally similar compounds-copper, iron, cobalt, zinc, and aluminum tetrasulfonated phthalocyanines (121). Structures a and d arise by simple electron transfer and only differ as to which is the lower lying energy level, whereas structures b and c require bond breaking reactions both in the formation and subsequent reaction of

420

PHILIP OEOROE

the complex. Further experiments are required to enable a choice t o be made. Preliminary measurements of the paramagnetic susceptibility show the iron in the metmyoglobin complex to be essentially ionjcally bound with p = 5.10 & 0.05 Bohr magnetons, close to the theoretical value of 4.90 calculated on a “spin only” basis for four unpaired electrons (George and Irvine, 121). This would be in accord with ionic structures of type a or b provided there was no large orbital contribution to the magnetic moment. I n the case of the secondary horseradish peroxidase complex, Theorell showed that the bonding is essentially covalent with xtn5 4,800 X c.g.s. units and it has been generally accepted that it is a covalent ferric compound with one unpaired electron. This value is somewhat uncertain since the solution examined contained in addition to the secondary complex some of the primary complex. More recently Theorell, Ehrenberg and Chance (122) report a better value of 3,500 X c.g.8. units for the methyl hydroperoxide secondary complex. This value is close to the susc.g.s. units, ceptibility of the peroxidase-cyanide complex, 2,970 X and they conclude by analogy that the complex containsi ferric iron essentially covalently bound like the cyanide complex. A susceptibility value of this magnitude is very high for a compound with only one unpaired electron where the contribution arising from electron spin is 1,270 X c.g.s. units at 20°C. Contributions from orbital moment are believed to account for such deviations observed generally with covalent ferric complexes of the hemoproteins (Pauling, 4). It is interesting however that the value of 3,500 X c.g.s. units for the complex is even a little greater than the theoretical value of 3,390 X 1 0 P c.g.8. units required for two unpaired electrons calculated on a “spin only” basis. This is the number required for a covalent hexacoordinated compound of quadrivalent iron or its derivatives. Since there is no a priori reason for expecting the orbital contribution to be large or small for the secondary complex the susceptibility value quoted does not exclude a quadrivalent iron type of structure. b. Oxidative Reactions of Hemoglobin and Myoglobin Involving Oxygen. There is no evidence that the autoxidation of these compounds is complicated by the participation of complexes of unknown structure and since the overall chemical change occurring is the simple oxidation of a ferrous compound to a ferric compound, these reactions are more suitable for kinetic analysis. The intramolecular reaction mechanism for hemoglobin oxidation proposed by Lemberg and Legge in which only the intermediate Hb4(02)z undergoes a reaction has been discussed above. The fact that the mechanism cannot account for the oxidation at high oxygen pressures makes it

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINR

421

extremely unlikely that it is correct. It is, however, possibIe t o extend this mechanism and permit reaction between free heme groups and

heme groups combined with oxygen in the intermediates Hb4(02) and Hba(O2)~as well as in Hb4(02)2, discarding the idea that in addition a hydrogen donor AH2 is necessary. Such a mechanism with “reacting pairs” of heme groups fits the experimental data far better than that suggested by Lemberg and Legge, and with the additional assumption that the pairs in Hb4(02)3 react a little faster than the pairs in the other intermediates a very close fit can be obtained. Thus it is not possible to reject this type of intramolecular mechanism outright, but there remain the inherent difficulties in understanding how such a reaction can take place between separated groups and why an identical reaction occurring intermolecularly and favored by collision should be excluded. The experimental studies of George and Stratmann (37) showing that the kinetics of myoglobin autoxidation are similar in every respect to those of hemoglobin suggest very strongly that a common reaction mechanism is at work. Since myoglobin contains only one heme group this would exclude all intramolecular mechanisms and focus attention on possible intermolec,ular mechanisms. A rate dquation for the autoxidation appropriate to an intermolecular mechanism which agrees well with Brooks’ experimental data has been discussed above. It has the form:

&- k a ( a - 2) . (1 - a ) ( a - ). dt

a--2

where (a - x) and x are the concentrations of unoxidized and oxidized hemoglobin or myoglobin and a the fraction of unoxygenated hemoglobin or myoglobin. The numerator in this equation is reminiscent of the corresponding term in the rate equation derived from the free radical mechanism for autoxidation referred to earlier, though this simple mechanism cannot apply in this case because the denominators in the rate equations are different, which means in effect that the free radical reaction would not be first order in unoxidized compound. It is possible to extend the free radical mechanism so that the derived rate equation has the required first order dependence by assuming that there is an auxiliary electron accepting group in hemoglobin and myoglobin that can act as cupric ions do in the ionic iron-hydrogen peroxide system catalyzing the reaction between Fe3+ and 0 2 - . To illustrate this the following symbols are necessary: 0-Fe2+, 0-Fe2+.02, and 0-Fe3+ represent hemoglobin, oxyhemoglobin, and methemoglobin with the electron accepting group in its oxidized state.

422

PHILIP GEORGE

When it accepts an electron, it is represented by @--Fez+ or e)-Fe2+*0z. The reaction scheme is then: 0-Fe2+.0z

{ (

O-FeZ+ 0-Fe2+

1 3

O-Fe3+

+

02-

+ HOz 2 O-Fe3+ + OZH-

O-Fez+Oz @-Fez+ @-Fe2+02

--*

+

4 0 2 - 4

+ 0-~e3+

{

@-Fez+

+ 0 2

@-Fe*+ 5 O-Fez+ .-+ fast

(

O-Fe2+.0z

+ 0--Fez+

followed by the rapid reaction of HzOz with two more hemoglobin molecules. The solution of the stationary state equations written in Brooks’ terminology is :

where kl, kz, ka, and k d are velocity constants of steps 1, 2, 3, and 4, respectively. If steps 4 and 5 predominate, i.e., the catalysis of step 3 which is the back reaction regenerating hemoglobin, then Eq. (2) reduces to which is identical in kinetic form with Eq. (1). The fate of the HzOz is not kinetically significant, if it reacts with the ferrous iron of heme the number 4 appears in Eq. (3) ; if it reacts entirely with other groups the number 4 is replaced by 2: it merely affects the stoichiometry, not the kinetics. Essentially this mechanism is a competition for HO2 radicals; only a fraction of those produced in the initial electron transfer from the ferrous iron of heme to 0 2 lead to an overall oxidation because of the back reaction regenerating heme. The predominant back reaction is that catalyzed by the electron-accepting group in the hemoglobin, and it is by this reaction that the iron is protected against oxidation. It is held in the ferrous state, not in a static sense by its structural environment, but in the dynamic sense that the primary step giving rnethemoglobin is obscured by catalyzed back reactions regenerating heme, thus giving a slow net oxidation. There is another type of reaction of hemoglobin which can readily be explained by a similar radical mechanism. Acid or pyridine denature Hbh(02)k and liberate two of the four 0 2 molecules, a third one being evidently required for the oxidation of four iron equivalents, while the

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

423

fourth is used in unspecified oxidation reactions. Lemberg and Legge (5) have taken this as evidence for the existence of the two XH2 groups per molecule and represent the reaction : Hb4(Oz)4

+ O 2 + 2H20 2HzX +

-+

0

2

+

Hb4* 402 Hb4*(OH)4i.e., oxidation of the Fe 4 2Hz0 2X

-+

Hb4*

+

Hbr(On)r(HzX)z+ Hba*(OH)r(X)z

+ 202

where the asterisk indicates denatured globin. Lemberg, Legge, and Lockwood (41) and later Lemberg (40) studied the oxidizing action of acidified Hb4(02)4 on ascorbic acid and biliverdin. In the case of ascorbic acid, oxidation occurs instantaneously on acidification if the ascorbic acid is already mixed with the Hb4(02)4, but far less is oxidized if it is added shortly after acidification. With biliverdin there was little evidence for the instantaneous oxidation, but a slow oxidation was observed which was attributed to the formation of H102 in the initial reaction. The results were explained by assuming the formation of active oxygen which reacted directly in an oxyhemoglobin-ascorbic acid complex. These observations may also be interpreted by the free radical mechanism, 02- again being formed as in step 1 of the autoxidation mechanism 402-

+ 4H+-+ 2HzOz + 202

where MetHb4* signifies the acid hematin protein complex. 0 2 - reacts rapidly with water giving HzO2 and 0 2 (George, 86), and this will predominate, for even if some acid hematin-denatured protein complex reacts with 02-the corresponding heme complex will autoxidize regenerating 02-. The oxidizing action of the system will depend on whether the substrate can react with 0 2 - or the HOz radical with which it is in ionic equilibrium. If it does not, oxidation can still occur through the peroxidatic action of the H 2 0 2and acid hematin-denatured protein complex. Ascorbic acid unlike biliverdin apparently reacts with 0 2 - . The fact that 02-(H02) is not an extremely powerful oxidizing agent has been shown in the recent experiments of Barb, Baxendale, George, and Hargrave (83). The action of pyridine on Hb4(02)4 can be explained in a similar way for it denatures the protein and parahematin derivatives are formed.

+

Hbr(O2)r 402-

pyridine

Globin and pyridine parahematin

4H+ + 2Hz02

+ 202

+ 402-

424

PHILIP GEORGE

These mechanisms explain immediately why usually only 50% of the total combined Oz appears as 0 2 gas. The fact is often expressed by saying 50% rather than the theoretical 75% is liberated. On this new interpretation the theoretical amount would be SO%,, and the actual amount obtained might be less if 0 2 - reacted with a substrate or more if the HzOZ formed were catalytically decomposed. An advantage of these mechanisms is that the stoichiometry is explained without the assumption of two XH2 groups. Other reactions of hemoglobin also permit a free radical interpretation, notably the coupled oxidation with ascorbic acid by molecular 0 2 which yields choleglobin, but further discussion requires a full kinetic analysis. Even though the denaturation reactions described above have not been examined kinetically it is worth emphasizing that their chief features can be explained by the formation of 0 2 - as in the mechanism advanced for the autoxidation. The liberation of an activated lo2molecule is no longer required-02- is the active oxygen.

VII. SUMMARY In comparing the reactions of hemoglobin, myoglobin, peroxidase, and catalase with molecular oxygen or hydrogen peroxide in relation t o the similar reactions of ionic iron the following conclusions can be drawn: 1. Many experimental studies now suggest a free radical mechanism in the case of ionic iron, but the essential nature of this evidence is indirect. Most important are the kinetic and thermochemical data. 2. Consideration of the possibility of similar free radical reactions occurring with the hemoproteins according to this thermochemical data suggests that the ferrous compounds should react as rapidly or more rapidly than free ferrous ions. The ferric compounds reacting with HOz radical or 0 2 - should still be rapid reactions but in the case of hydrogen peroxide or OzH- the reaction should proceed extremely slowly. 3. Kinetic studies of the autoxidation of hemoglobin and myoglobin do not favor intramolecular mechanisms involving reaction between separated groups on the protein molecule, but the form of the rate equations resembles that found when two valency states of a metal ion compete for the same radical intermediate. A free radical mechanism based on competition for the HOz radical can be developed to account for the observed results. 4. It is the reaction of the ferric forms of these hemoproteins with hydrogen peroxide that provides the sharpest contra,& to ionic iron. Complexes are formed in both systems. Although ferric ion only gives an ion pair complex Fe02H2+,whereas peroxidase can give three com-

THE SPECIFIC REACTIONS OF IRON IN SOME HEMOPROTEINS

425

plexes, depending on experimental conditions, the first of which probably has this ion pair structure, the contrast is in the type of reaction undergone by the complex. The ferric ion complex plays no independent kinetic role, the dominant reaction being electron transfer resulting in the partial oxidation of the peroxide molecule: in the hemoprotein peroxide complexes the peroxide is utilized in reactions involving the reduction of peroxide t o water. So little is known about the structure of these complexes t h a t a detailed discussion of reaction mechanisms must await further experimental evidence. I n the reaction of ionic iron there is some indication that a “ferryl” ion, Fe02+ may take part under certain experimental conditions. A compound with this structure would adequately explain part of the known behavior of the secondary complexes with peroxidase and catalase and the single complex formed with methemogIobin or metmyoglobin. I n the light of these conclusions and the data on which they are based, the effect of coordination within the porphyrin ring and t o the protein molecule on the catalytic action of the iron atom is potentially fourfold. 1. This coordination makes the complexes water soluble over a large pH range. This effect is important in itself for if ionic iron were not precipitated as the hydroxides, a t p H 10.46 i t would have a catalytic activity comparable t o that of catalase in the straightforward decomposition of hydrogen peroxide. 2. The ionization potential of the ferrous compounds in aqueous solution is lowered by the coordination, which has the effect of making any endothermic reaction of free ferrous ion less endothermic and hence potentially faster in the case of the hemoproteins. 3. Conversely the reaction of the ferric compounds will be potentially less exothermic, or endothermic and hence slower. I n this respect the coordination of the iron atom appears t o make possible a new reaction path in the case of hydrogen peroxide and alkyl hydroperoxides involving at first complex formation with the peroxide molecule and then the production of further complexes of unknown structure but which function as electron or hydrogen acceptors. 4. Another possible result of the coordination is t o allow electron accepting or donating groups on the protein molecule t o enter into competition reactions which could lead t o the predominance of certain reactions resulting in a selective or specific overall reaction.

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