Chemistry of iron and copper in radical reactions

C.A. Rice-Evans and R.H. Burdon (Eds.), Free Radical Damage and its Control 6 1994 Elsevier Science B.V All rights reserved

3

CHAPTER 1

Chemistry of iron and copper in radical reactions W.H. KOPPENOL Departments of Chemistry and Biochemistq, and Biodynamics Institute, Louisiana State University, Baton Rouge, LA 70803, USA

Abbreviations adp amp atp dtpa edda edta

adenosinediphosphate adenosinemonophosphate adenosinetriphosphate diethylenetriamine-N, N, M,M'.MIpentaacetate ethylenediamine-N, A"-diacetate ethylenediamine-N, N, M ,M-tetraacetate

gtp hedta nta phen PQ'+ utp

guanosine triphosphate (N-hydroxyethy1)ethylenediamineN, M,M-triacetate nitrilotriacetate 1,lO-phenanthroline paraquat radical uridine-5'-triphosphate

I . Introduction In a comparison of several elements it was shown by George[1] in 1965 that oxygen is unique because its reduction by organic compounds is favourable and because the reaction product, water, is not toxic. As such oxygen is the best element on which to base life. Yet oxygen also plays an important role in free radical biology in that it is also essential in the initiation of, and greatly amplifies, damage to biomolecules. Oxyradicals have been implicated in numerous diseases and disorders. Iron and copper catalyse the formation of oxyradicals. Three reactions are relevant in this context: (1) Autoxidation of metal complexes may yield the superoxide radical which by itself is not very reactive, but is a precursor of more reactive radical species. (2) The one-electron reduction of hydrogen peroxide the Fenton reaction - results in hydroxyl radicals via a higher oxidation state of iron [2]. (3) A similar reaction with organic peroxides leads to alkoxyl radicals, although a recent report alleges that hydroxyl radicals are also formed [3]. There is a fourth radical, the formation of which does not require mediation by a metal complex. This is the alkyldioxyl radical, ROO', which is formed at a

A

TABLE 1 Reduction potentials of oxyradicals” Inorganic Couple

Organic E“‘W 7 ) (V)

Couple

One-electron reduction potentials 0210; -0.33

EO’(PH7) (V)

~

HOiIH202

1.07

ROO’IROOH

1.o

‘OHIH20 H202I’OH, H2O

2.3 1 0.32

RO‘IROH ROOHIRO’, H20 bis-allylic’lbis-allylicH

1.7 1.9 0.6

Two-electron reduction potential H20212H20

1.32

ROOHIROH, H20

1.8

”

Data for H- and 0-containing radicals were taken from a recent compilation [4]. The values for the carbon containing radicals arc estimates derived from bond-dissociation enthalpies [5]. The value for E”’(ROO’IRO0H) has recently been confirmed experimentally [6].

nearly diffusion-controlled rate from an alkyl radical and dioxygen. As shown in Table 1, the hydroxyl, the alkoxyl and the alkyldioxyl radical are oxidizing species. For a quantitative description one needs to know rate constants for the formation of these radicals and for their subsequent reactions in order to determine which reaction will dominate under physiological conditions. This requires inter aha knowledge of the precise chemical composition of the cell; we are still a long way from this goal. In principle, there are three mechanism for damage: a single reaction, a chain reaction, or a branching mechanism. A single reaction is not likely to lead to extensive damage. However, when a radical like the hydroxyl radical reacts with a biomolecule, another radical is created. A chain reaction ensues that stops only when a radical reacts with another radical or with a transition-metal ion. For extensive damage to occur it might be necessary that branching occurs. For instance, superoxide production may lead to lipid peroxidation; alkenals formed as products of that process are substrates for xanthine oxidase [7]; more superoxide is produced, and a new chain reaction is started. Similarly, during ischaemia atp is converted to hypoxanthine [8]. Any iron that was tightly bound to atp is now bound elsewhere, possibly in a more “open” complex. Since rate constants increase with a decrease in coordination number (see below), such an iron complex is likely to be more reactive. Given the concentration of “reaction sites” in vivo and the magnitude of the relevant rate constants it is not possible to intercept effectively the hydroxyl, alkoxyl or alkyldioxyl radicals [9]. For less reactive compounds,

5

such as hydrogen peroxide and superoxide, nature has developed enzymes to dispose of them. The strategy adopted by nature is threefold: (1) interception with superoxide dismutase and proteins such as catalase and glutathione peroxidase that react with hydrogen peroxide and small alkyl hydroperoxides; (2) repair of water-soluble biomolecules with glutathione, and ( 3 ) inhibition of lipid peroxidation with vitaminE, which might [lo], or might not [ l l ] , be regenerated by vitamin C. Part of the defence mechanism may be that radicals are interconverted to superoxide via the glutathione radical; superoxide, acting as a radical sink, is subsequently scavenged by superoxide dismutase [12]. The overall energetics of these reactions are extremely favourable [ 131. Excess amounts of transition metals, in particular iron and copper, are toxic. For instance, it has recently been suggested that excess iron plays a role in heart disease [ 14,151. Transition-metal ions are sequestered by proteins: iron in ferritin and transferrins [ 161, and copper in caeruloplasmin. However, a small concentration of low-molecular-weight complexes is likely to be present at all times because of transfer of metals from storage proteins to metalloproteins, and from the turnover of these proteins. Under oxidative stress this pool of iron is increased due to reductive mobilisation and destruction of ferritin [ 17-25]. Desferrioxamine, being a stronger complexing agent than the naturally occurring ligands, chelates iron and prevents oxidative injury in hepatocytes [26]. The precise nature of the low-molecular-weight complexes present in viva is not known with certainty. Evidence has been presented that iron bound to atp, amp [27], gtp [28] and possibly citrate [29] may be present in tissues in the micromolar range [21]. No information is available on copper. Some metal- (especially copper) complexes catalyse the dismutation of superoxide at rates that compare favourably with catalysis by superoxide dismutase. One could therefore argue that the presence of such complexes in viva might be beneficial. There are, however, additional considerations: (1) such metal complexes may also reduce hydrogen peroxide, which could result in the formation of hydroxyl radicals, and (2) it is extremely likely that the metal will be displaced from its ligands (even when those ligands are present in excess), and becomes bound to a biomolecule, thereby becoming less active as a superoxide dismutase mimic. As an example, copper binds well to DNA and catalyses the formation of hydroxyl radicals in the presence of hydrogen peroxide and ascorbate [30]. Both the reduction of superoxide and that of hydrogen peroxide appear to be inner-sphere reactions; that is, a ligand of the metal ion has to be replaced by superoxide or hydrogen peroxide for the reaction to take place. For superoxide this involves overlap between a metal d-orbital and its own accessible T * orbital. Reduction of hydrogen peroxide involves electron transfer to an empty 0 * orbital which is not very accessible [31]. Thus, reductions of hydrogen peroxide are generally slower than those of superoxide. The reductions of alkylhydroperoxides are even slower, due to steric hindrance [32,33].

6

This review is concerned with the quantitative aspects of metal-catalysed oxyradical reactions. As such one will find discussions of structures of metal complexes, rate constants and reduction potentials, not unlike our review of 1985 [34]. Two areas related to the role of transition metals in radical chemistry and biology have been reviewed recently; these are the metal-ion-catalysed oxidation of proteins [35] and the role of iron in oxygen-mediated toxicities [36]. These topics will not be discussed in detail in this review. Related to this work is a review on the role of transition metals in autoxidation reactions [37]. Additional information can be obtained from Afanas'ev's two volumes on superoxide [38,39]. This subject is also treated in a more general and less quantitative manner by Halliwell and Gutteridge [40].

2. Autoxidation reactions 2.1. Oxygen

It is well known that oxygen does not react directly with organic molecules because of spin restrictions: ground-state oxygen is a triplet molecule, and most organic molecules are in the singlet state (see ref. [37]). In the past we have explained this phenomenon qualitatively in the following fashion: Prior to a reaction, overlap is necessary between orbitals of the reactants; this can only occur rapidly between half-filled orbitals (of the proper symmetry) which organic molecules generally do not have [3 11. Similarly, triplet oxygen and organic radicals react at near difhsion-limited rates because both have half-filled orbitals. Often, authors of reviews try to explain the singlet and triplet states of oxygen with a diagram as depicted in the middle column of Fig. 1 [36,40,41]. Such primitive representations have been criticised [42] because, for one, it does not show that there are 6 ways in which 2 electrons can be distributed over two orbitals. There are 6 different microstates that belong to three different energy levels: ground state (3C,) oxygen is three-fold degenerate, the 'A,state is two-fold degenerate, and the 'CH state is not degenerate. A correct orbital occupation-energy diagram, taken from refs. [42,43], is depicted in the righthand column of Fig. 1. In the lowest energy state, 'Xi,the two electrons move in mutually perpendicular planes, minimizing repulsion, with parallel spin. In the highest state, 'Xi,which because of its extremely short lifetime is biologically not relevant, the electrons move in the same plane with paired spins, while in the 'A, state both electron orientations occur. Thus, this state can undergo both two-point additions and single-point attachments [42,43]. A well-known twopoint reaction is the addition of A, to a double bond. Single point attachments do not lead to a reaction with singlet organic molecules, but would allow 'A, to react like 3X; oxygen with radical species.

'

State

Orbital Assiament Primitive

’cg+ Ox@, 1

A,

oxoy

7

Correct

oxoy oxoy +

oxoy - 0, oy @,ay - ox0,

Fig. 1. Spin-orbital diagram of the different states of oxygen. The x and y refer to the two perpendicular antibonding orbitals of oxygen. On the right this diagram depicts the real wave-functions for the lowest electronic states. From ref. [43].

-A

2.2. Thermodynamics

At low pH the iron(I1) ion is stable with respect to oxidation, due to the high value of the reduction potential of the Fe3’/Fe2+ couple, 0.77V versus the normal hydrogen electrode. Above pH 2.1, Fe(III), but not Fe(II), hydrolyses, which results in a reduction potential that decreases with 59 mV per pH unit to a value of 0.48V at pH7. This applies only to very dilute solutions, since iron(II1) hydroxide precipitates above pH 3. Complexation by aminopolycarboxylates, such as edta, which provide mainly oxygen as donor atoms, also reduces the reduction potential, generally to a value near 0.1 V [34]. The standard reduction potential of the oxygen/superoxide couple is -0.33V (see Fig. 2), independent of pH[44,45]. Although in such an instance the one-electron reduction of oxygen by such metal complexes is thermodynamically unfavourable by approximately 10kcal/mol, the reaction proceeds because the product, superoxide, disappears by disproportionation. In contrast, reduction by two electrons to hydrogen peroxide is favourable: the Gibbs energy change is -8.1 kcal per two moles of Fe(I1) edta, as calculated from E0’(02/H202) = 0.305 V at pH 7 [44] and the reduction potential of the Fe(II1)-/Fe(II)-edta couple of 0.12 V [46]. It has been suggested that a (dioxygen)iron(II), or “perferryl” I complex, a likely intermediate in the autoxidation of iron(II), could abstract an allylic hydrogen and initiate lipid peroxidation [48]. Such complexes are weak oxidants at best, as has been shown before[49] and, with the exception of iron(I1) edta [50], have not been observed. Constraints on the reduction potential

’

The name “perferryl”, indicating an oxidation state beyond that of ferry], iron(IV), is not recommended by the current IUPAC guidelines for the nomenclature of inorganic chemistry [47]. This name would only be defensible if both oxygen were attached to the iron, which they are not. The use of this misleading name should be discontinued.

8

!

pH 7 po,= 1 atm

HO’

t

03.

0 0

w c -1

-2 -2

0

-1

n

Fig. 2. Oxidation state diagram of oxygen at pH 7 at otherwise standard conditions ( 1 molal concentrations, 1 atm for gases). The x-axis gives the oxidation state, the y-axis the product of reduction potential and oxidation state. As such the slope represents the reduction potential. Adapted from ref. [4]. A compound that lies above a line joining its neighbours is unstable with respect to disproportionation, as is the case for superoxide and hydrogen peroxide. The line from hydrogen peroxide to the middle of the water-hydroxyl line represents the one-electron reduction potential of the couple H202I’OH, H 2 0 .

Eo’(HLFe1102/HLFe111, H202) come from the following thermodynamic cycle and considerations. If such a complex were to be an initiator of oxyradical damage, one might expect that approximately 1YOof the low-molecular-weight iron(I1) be complexed to oxygen at a cellular oxygen tension of, say, 0.01 atm. This requires a standard Gibbs energy change of Okcal/mol. In the following sequence of reactions HL represents a ligand with a covalently bound hydrogen: HLFe(II)02 0 2

+ HLFe(I1) + 0 2

+ 2H’ + 2e-

-+

H202

HLFe(I1) -+ HLFe(II1) + eHLFe(II)02 + 2Hf + e-

-+

(AGO’ = 0 kcaVmol),

(1)

(Lo’= 0.305 V),

(2)

-0.1 V),

(3)

(EO’

=

HLFe(II1) + H202

(EO’

=

0.2 V).

(4)

The reduction potential of 0.2V for Reaction (4) at pH7 depends very much on the reduction potential of the Fe(III)/Fe(II) couple, Reaction ( 3 ) . The 0.1 V assumed here for that half-reaction is close to that of various ironaminopolycarboxylate complexes. The uncertainty in our reduction potential for Reaction (4) is estimated at 0.2V The abstraction of a doubly allylic hydrogen is estimated to require a reduction potential of 0.6V (see Table l),

9

which makes the reaction with a (dioxygen)iron(II) complex unfavourable by approximately 10 kcal/mol. A variation on the (dioxygen)iron(II) complex, an Fel102FeII1intermediate, was proposed by Aust and coworkers as the instigator of oxyradical damage [37,51]. There is no thermodynamic data available that allows one to calculate how oxidising such a complex would be. It is conceivable that an equal mixture of iron(I1) and iron(II1) compounds imposes a reduction potential on the system that is favourable for catalysis of lipid peroxidation. Not many reduction potentials are known for copper complexes. That of the Cu2+/Cui couple is 0.16 V. Since E"(Cu+lCu") is 0.52 V, the disproportionation of Cu' to Cuo and Cu2+ is favourable. This reaction does indeed occur, which makes is impossible to study stable copper(1) solutions. Reduction potentials of copper(I1)-/copper(I)-( 1,l O-phenanthro1ine)z and a few derivatives have been calculated from a kinetic analysis of appropriate rate constants: values range from 108 mV for the 5-methyl-1,lO-phenanthroline complex to 219mV for the complex with a nitro group at the 5 position [52]. Values of 0.17V and 0.12V are given by Phillips and Williams [53] for the phenanthroline and bipyridine complexes, respectively. Such complexes can thermodynamically catalyse both the superoxide dismutation and the one-electron reduction of hydrogen peroxide (see below). 2.3. Kinetics and mechanisms

The rate of autoxidation can be calculated if the Gibbs energy of reaction and the rate constant for the reverse reaction, the reduction of the metal complex by superoxide, are known. The reduction potential of the Fe(II1)-/Fe(II)-edta complex is 0.12V at pH7[46], and that of the oxygen/superoxide couple is -0.33V (see above), which results in a ArxnGo'of -10.6kcal/mole at pH7 for the reduction of Fe(I1I) edta by superoxide. This reaction proceeds with a rate constant of -5 x lo6 M-' s-l [54,55]. A rate constant of -30M-Is-' for the oxidation of iron(I1) edta by oxygen at pH7 is calculated from the Gibbs energy change and the rate constant. It follows from the pH dependence of the reaction of iron(II1) edta with superoxide [55] that the autoxidation would be faster at lower pH. Near neutral pH, rate constants of 600M-' s-' [56] and 270M-' s-' [57] have been reported, much faster than the 30M-' s-' estimated above. We observed that the autoxidation of dilute (micromolar) iron(I1) edta solutions is first-order in iron(I1) complex and in oxygen, with a rate constant of 110M-' s-' at neutral pH, in better agreement with the thermodynamically predicted value (Koppenol and Rush, unpublished). The first study of the kinetics of autoxidation reactions of a number of iron(I1) salts was published in 1901 by McBain[58]. It established that the reaction is first order in oxygen pressure and second order in iron(I1). In contrast,

10

a similar study on iron(II)hydrogencarbonate, published in 1907 by Just [59] showed that the reaction is first order in iron(I1) and in oxygen. This work is also noteworthy for another reason, for it mentions for the first time the superoxide anion. Most other studies report a second-order dependence in Fe(I1) [60-661. A reaction mechanism was proposed by Weiss[67] in 1935 which covers Just's observations [59]. It involves first the formation of superoxide, which after protonation oxidises another Fe(I1): Fe2++ 0

+ Oi-,

2 + Fe3+

(5)

+ Hf -+ HO;, Fe2++ HO; + Fe3++ HOT.

(6)

0;-

(7)

These reactions are followed by the reduction of hydrogen peroxide which consumes two more iron(I1) ions. According to Reaction (9,the autoxidation of iron would be first-order in iron(I1). However, for most other anions the autoxidations are second-order in iron(II), and for that reason this mechanism was criticised by George [62]. He studied the oxidation of iron(I1)perchlorate solutions at higher oxygen pressures and showed that this reaction is secondorder in iron(I1) and first-order in oxygen [62]. His mechanism involves formation of an intermediate dioxygen-iron(I1) complex, Reaction (8), followed by rate-limiting reaction with a second iron(II), Reaction (9): Fe2++ 0

2 + Fe2+02,

Fe2+02+ Fe2++ 2H'

(8)

+ H202.

+ 2Fe3+

(9) The studies mentioned above were carried out at lowpH. More relevant to this discussion are studies at neutral pH in the presence of chelating agents. At this pH the hydrodioxyl radical is present as superoxide, which is less likely to oxidise iron(I1) complexes, although in most instances this would be thermodynamically feasible, E0'(O;-/H202) being 0.94 V at pH 7 [44]. Such an oxidation would have to be fast to compete with the dismutation reaction. The oxidation of iron(I1) by dioxygen is pH dependent: at pH7.03 the halflife of this process, 2700s, is -10 times greater than that at pH7.45 [68]. The presence of chelating agents drastically decreases the halflife. In the case of edta the halflife is 10s near neutral pH. No rate constants were reported. The autoxidations of iron(I1) aminopolycarboxylates proceed with rate constants of 270, 100, 80 and 7M-' s-' for the ligands edta, hedta, nta and dtpa, respectively [57], and a relationship between the rate constant for the oxidation and the ratio of the stability constants of the iron(II1) chelate to that of iron(I1) was established. Similar autoxidation rates were obtained by another group [69]. An intermediate iron(II)edta-oxygen complex has been postulated [50]. This complex is believed to become protonated at lower pH to

I1

form the hydrodioxyl radical, as proposed for unchelated iron [67], and reacts at higher pH's with excess iron(I1)edta to form hydrogen peroxide. The protein apotransferrin was also shown to increase the rate of autoxidation [70], probably by removing iron(II1) from equilibrium. At pH 7.0, phosphate increases the rate of oxygen consumption in iron(I1)-solutions, and this process is slowed by dtpa [71]. Recently the curious observation was made that phosphate slows the rate of autoxidation, although no rate constants were given[72]. It can be concluded from this overview that the autoxidation of physiologically relevant iron(I1) complexes has not been well characterised, and that conflicting reports exist about the effect of phosphate on the autoxidation reaction. The reverse reaction, the reduction of iron(II1) complexes by superoxide, proceeds with rate constants varying from 1.9x 1O6 M - s- for edta to 7.6x 1 0 5 s-~ for ~ hedta, to negligible for dtpa[54,55,73]. The same trend as seen for the reaction of iron(I1) complexes with hydrogen peroxide (see below) is observed here: The rate constant decreases when the number of ligand atoms provided by the chelating agent increases. Superoxide forms an adduct with iron(I1) complexes [54,56]; this complex is also formed from the iron(II1) complex with hydrogen peroxide [74]. Recently it was reported that iron(II1)citrate undergoes autoreduction [75]. This process is known to be photocatalysed and was described more than 50 years ago [60]. The reduction of oxygen by copper(1) is faster than that of the iron(I1) complexes: 5 x 1o4M-' s- for Cu'@hen)2 [52] and 4 x 104M-' s- for Cu'(histidine)2 [76]. It is this relatively fast autoxidation that limits the usefulness of copper complexes as mimics of superoxide dismutase under conditions of high superoxide concentrations [77]. Copper(I1) catalyses the dismutation of superoxide at near diffusion-controlled rates: k,,, = 8 x lo9M-' s-' [78,79].

'

'

3. Fenton reactions 3.1. Introduction

The reaction of iron(I1) with hydrogen peroxide is named after H.J.H. Fenton who observed in 1876 [80] that addition of hydrogen peroxide to a mixture of tartaric acid and iron(I1) sulfate, followed by addition of base, resulted in a dark purple colour. A full account was published in 1894 [81]. Having first used it as an analytical tool for the assay of tartaric acid, Fenton then employed this type of reaction to study the oxidation of a variety of organic compounds. The Fenton reaction is generally considered to yield the hydroxyl radical [82], as follows from spin-trapping [83,84] and hydroxylation [85-891 studies. The hydroxyl radical is a strongly oxidizing agent, E"('OWH20) = 2.73 V [90], and attacks various small molecules with rates of 10s-lO'oM-'s-l, while its

12

reaction with proteins is diffusion-controlled [9 11. In the presence of oxygen the peroxyl radical is formed [92], which can start various chain reactions [93]. The propagation reactions are well understood and, as mentioned above, are responsible for far more damage than the initiating event. While the concept of the hydroxyl radical as an initiator has received wide support, recent evidence suggests that a higher oxidation state of iron might be involved. This concept is not new: as early as 1932 it was proposed that a higher oxidation state of iron, the ferry12 ion (Fe02'), might be involved in the decomposition of hydrogen peroxide [94]. This concept formed an integral part of the chain reaction proposed by Cahill and Taube in 1952 [95]. Currently there seems to be a consensus that at lowpH the hydroxyl radical is formed [82,96]. However, the situation is more complex at neutral pH when iron is present in chelated form. The failure of common hydroxyl-radical scavengers to inhibit, for instance, the formation of ethylene from methionine, resulted in the postulation of electron-donor or "crypto-'OH' complexes [97,98]. Such a complex or ferryl compound should be fairly oxidizing to show more or less the same reactivity as the hydroxyl radical. For instance, a reduction potential of 1.2V is required to abstract an a-hydrogen from methanol [99], and such a value would seem to be a lower limit for a ferryl species. A thermodynamic derivation [IOO] suggests a value in excess of 0.9 V for a hypothetical Fe(1V)Ee(II1)-edta couple. However, little is known directly about the structure and reactivity of high-valent iron in aqueous solution, with the exception of that of ferryl porphyrins [ 102,1031, and some spectroscopic information and decay kinetics of ferryl and ferrate, FeOi-, in alkaline solution [ 104,1051. Ferry1 may be represented as Fe"(H202), [Fe'v=0]2+, Fe"'-O-, or FeIV(OH-)2. 3.2. Thermodynamics The one-electron reduction of hydrogen peroxide is thermodynamically favourable if the reduction potential of the metal complex is 0.32V or less at pH7 [4]. This value is based on the reduction potential of the HO'/H20 couple, according to the following equation: E0'(H202/HO', H20) = 2E0'(H202/H20)

- E0'(HO'/H20).

(10)

The reduction potential for the hydroxyl/water couple was not precisely known until recently [ 106,107]. Based on older values for the hydroxyl radical/water couple one will find higher values of 0.8 V [ 1081 or 0.46 V [34] in earlier papers by the present author. The implication of the value of 0.32V for the reduction The term "ferryl" is commonly used to describe an oxidising iron species derived from the reaction of hydrogen peroxide by iron(I1) complexes, although the expression oxoiron(1V) complex is better[lOl].

13

of hydrogen peroxide is that complexes such as tris-1,lO-phenanthrolineiron(I1) and tris-2,2’-bipyridine iron(I1) are unlikely to reduce hydrogen peroxide since the reduction potentials of the respective iron complexes are in excess of 1 V, which makes the reaction unfavourable by 16 kcal/mol.

3.3. Kinetics Rate constants have been determined for the reduction of hydrogen peroxide by iron(I1) and a number of iron(I1) complexes. These rate constants have been compiled in Table 2. It is immediately clear that there is not much agreement between the results of various groups. However, there is a discernable trend: metal complexes with more water-accessible coordination sites react faster. Graf et al. [ 1171 have commented upon the importance of coordinated water molecules for the Fenton reaction. It is also clear that the rate of the Fenton reaction for a chelated complex near neutral pH is much faster than that of aqueous iron(I1) at low pH. The use of the low-pH value of 76M-I s-’ in a recent calculation [ 1181 of the flux of hydroxyl radicals in a cell gives an estimate that is at least two orders of magnitude too low. The rate constants for the reduction of hydrogen peroxide by copper(1) phenanthroline and aqueous copper(1) are 1.1 x lo3 M-’ s-’ [119] and 4.1x103M-’ s-’ [120], respectively. 3.4. Intermediates

It has been argued that at neutral pH the Fenton reaction proceeds via an intermediate. Such an intermediate could be a oxoiron(1V) compound as shown in the following scheme:

>ieoH2+

I

+ HO*

- ’I

+

\je4”OH

-

Scheme 1

‘OH-

ferry1

I

site - specific damage

Loss of water from the Fe”(OH-)2 species would yield Fe02+, a compound formed in the oxidation of iron(II1) haems by hydrogen peroxide (as well as an oxidised porphyrin ligand). One might ask the question whether copper can form a cupryl, or oxocopper(III), species. It has been argued that this is not

14 TABLE 2 Rate constants for the Fenton reactiona Rate constant (k/103,M-ls-' ) reported byb:

Complex H

aqua'

41.5

dtpa

-

edta

-

hedta nta edda atp

BFA

SW

60

0.51

-

G

200 0.80d 13.5s

-

RR -

WNDF 75

YP -

1.37

-

0.4 1

10s

-

-

14

7.1h loh 9.7h

-

-

-

-

-

-

23h

-

-

-

-

-

-

-

-

-

-

4.9h

-

-

-

-

-

-

-

-

-

5.3h

-

-

-

2.7h

-

9.10 16.7 18.4 -

7 -

-

7.0d 42d 3Oe

-

7Se 6Af 1I f 5.2f 4.9' -

-

-

4J

-

utp

-

-

citrate phosphate pyrophosphate tartrate

-

-

-

-

-

-

-

-

-

-

-

a

K

lOOf -

-

8.2 20

Error limits are given in the original papers. Abbreviations and remarks: H, Hardwick [lo91 (20.2"C, 0.1 N , HC104); this reference contains a discussion of earlier work. BFA, Borggaard et al. [ 1101 (20"C, pH 6, 0.2 M ionic strength); values apply to unprotonated species and were determined polarographically. SW, Sutton and Winterbourn[ll I] (pH7.4, variable ionic strength). K, G, Various references (see notes below). RR, Rahhal and Richter [ 1121 (pH 7.0); rapid mixing study. WNDF, Wink et al. [113]; UV spectroscopy. YP, Yamazaki and Piette [I 141; ESR-flow study. Values for aquo are in k (M-' s-'). Rush and Koppenol[96] (25"C, pH 7.2, 38 mM ionic strength); stopped-flow study. Rush and Koppenol[2]; stopped-flow study. Rush et al. [32]; stopped-flow study. Gilbert and Jeff [ 1151, pH 4; ESR flow study. Croft et al. [I 161, pH 7; ESR flow study.

possible because formation of such a species requires the presence of an empty t2g metal orbital, such that a T-bond between one of these orbitals and a full 2p orbital of oxygen(2-) can be formed [121]. Copper(II1) still has 8 d-electrons, and there are no empty t2s orbitals. Thus, if a higher oxidation state of copper is involved, it is copper in the 3+ oxidation state, not a copperoxo(l+) species. However, iron(1V) has only 4 d-electrons, allowing the formation of one such .7r -metal-oxo bond.

15

At low pH the reaction of Fe(II)aq with hydrogen peroxide leads to the hydroxyl radical [82,96,122] and spectral evidence for an intermediate has recently been obtained [ 1131. At neutral pH, in the presence of chelating agents, the situation is more complicated, as reviewed in 1989 [123]. We have investigated four reactions that indicate that an intermediate is involved. The first is the oxidation of ferrocytochromec by a mixture of iron(I1) edta and hydrogen peroxide [ 1241. Had the hydroxyl radical been formed, cytochrome c would have been degraded, not just oxidised [125]. The interaction between ferric cytochrome c and the hydroxyl radical is peculiar in that a radical is created on the surface which reduces the haem [126]. The second and third are based on the use of scavengers. When a scavenger reacts with the hydroxyl radical the resulting radical is either reducing, oxidizing or neutral [ 1271. Thus, when hydrogen peroxide is mixed with excess iron(I1) chelate the absorbance change in the UV corresponds to zero Fe(II1) formed per hydrogen peroxide consumed (no absorbance change), 2 Fe(III)/H202, or 1 Fe(III)/H202, respectively. The scavenger tert-butanol reacts with the hydroxyl radical to form a radical that decays by reacting with another tert-butanol radical. If this is the case, then the reaction of an iron(I1) complex with hydrogen peroxide should yield one iron(II1) per hydrogen peroxide. However, we observed that tert-butanol is unable to prevent the oxidation of a second iron(I1) complex [96,128]. This can be explained by Reactions (1 1) and 12: H202 + HLFe2+ + HLFe"(H202), HLFe"(H202)

+ HLFe2++ 2H+ + 2HLFe3++ H2O.

(1 1) (12)

The experimental observation of 1.7-2.0 mol HLFe(I1) oxidised per mol of hydrogen peroxide in the presence of tert-butanol was confirmed by Rahhal and Richter [ 1121. They showed that the oxidizing species reacted equally fast with Fe2+-dtpaas with hydrogen peroxide, from which it was concluded that the oxidizing species could not be the hydroxyl radical. If the hydroxyl radical had been formed, the scavenger tert-butanol would have intercepted it and the oxidation of the second iron(I1) would have been prevented, because the tert-butanol radical is believed to be inactive. Gilbert and Jeff [ 1 151 offered an alternative explanation. They suggested that the hydroxyl radical is formed, and that the tert-butanol radical oxidises the iron(I1) chelate. This reaction is followed by a reductive elimination [ 1291. This is shown in Reactions (13)-( 15), in which the chelating agent HL is edta: H202 + HLFe2++ H+ + HO' HO'

+ HLFe3++ H20,

+ (CH3)3COH + 'CH2(CH3)2COH + H20,

'CH2(CH3)2OH + HLFe2++ H+ --+ CH2=C(CH3)2 + HLFe3++ H20.

(13) (14) (15)

Recently, evidence for this scheme was presented in that the tert-butanol radical was observed by flow-ESR; a rate constant of 2 x 106M-' s- for Reaction ( 15) has been proposed [116]. It seems likely now that tert-butanol cannot be used to distinguish between the hydroxyl radical and higher oxidation states. The third system is based on the scavenger formate. The dioxocarbonate( 1 -) radical formed, CO;-, is a strongly reducing radical, Eo(C02/CO;-) = - 1.8 V [99]. When generated from the Fenton reaction it is expected to reduce the metal and no absorbance change should result. However, we observed an intermediate with absorption maxima near 300 and 410nm that we ascribed to a compound with iron-carbon o-bonds [2,96,123], similar to those investigated by Cohen and Meyerstein [130]:

'

HLFe'"(OH-)Z + HCO,

+ 'LFe"

+ C0;- + H20 + Ht .

(16)

Intermediates are only observed for the ligands hedta, nta and edda and decay by second-order kinetics. Goldstein et al. [ 1311 have shown that C0;- reacts with iron(I1) nta and iron(I1) hedta to form an intermediate with a spectrum that, above 320 nm, is nearly identical to that described above. Our observations [2,96] could be explained [ 1311 by Reactions (1 3), (17) and (1 8): HzOz+ HLFe2' + H+ + HLFe3++ 'OH + H20,

(13)

'OH + HCO,

C0;- + H20,

(17)

+ HLFe2++ HLFe"-CO,.

(18)

CO;

-+

However, addition of hexaamminecobalt(III), which scavenges CO; -, reduced the transient absorbance change by half, but did not eliminate it[123]. We concluded that this absorbance may be caused by ligand radicals forming crbonds to iron@), which are not affected by hexaamminecobalt(III), and by C0;- bound to iron(I1). Evidence for a ligand radical, and therefore for a higher oxidation state of iron, also follows from the observation that formate does not protect the ligands edda and nta from being degraded to the extent that iron(II1) oxide precipitates from solution [2]. Using formate, Sutton and Winterbourn [ 1 1 1,132,1331 also presented evidence that the reaction of Fe2+-edta and aqueous, unchelated, iron(I1) with hydrogen peroxide involves in part a higher oxidation state of iron. This experiment involves a chain of reactions, which on the basis of the known rate constants, would have a very large number of cycles before it would come to an end. Instead, fewer than 10 cycles were observed, see below. These experiments have been extended to other ligands. In rapid-mix experiments a solution of iron(I1) complex [or iron(II)], formate and PQ' is mixed with hydrogen peroxide and the ratio of carbon dioxide produced per paraquat oxidised (the number of cycles) is determined. Alternatively, paraquat radicals are generated

17

continuously by ionizing radiation in the presence of formate, iron(II1) complex and hydrogen peroxide, and the same ratio is determined. The following chain reactions are believed to take place: PQ"

+ HLFe3+ + HLFe2++ PQ2',

H202 + HLFe2++ H+ -+ HLFe3++ 'OH 'OH

+ HCO,

C0;- + PQ2'

(19)

+ H20,

+ H20,

--+

C0;-

-+

C02 + PQ".

(13) (17)

Reactions (10) and (21) end the chain: H202 + HLFe2+

--f

HLFe"(H202),

HLFe"(H202) + PQ'

-+

HLFe3++ PQ2'

+ 20H-.

(10) (21)

The numbers of cycles are greater for the continuous irradiation experiments, but decrease from 12 for edta, to 9 for hedta, to 4 for edda, to nearly 3 for nta. The rapid-mix experiments give about half as many cycles as do the continuous-irradiation ones. Unchelated iron supports an even smaller number of cycles. Similar results were obtained when superoxide was produced by xanthine oxidase and hypoxanthine. In that system superoxide reduces the iron(II1) complex, which starts the chain. The numbers of cycles observed are: edta, 7.8; hedta, 6.5; edda, 1.7; and nta, 1.8 (Winterbourn, Sutton and Koppenol, unpublished). It would seem that Reaction (13) predominates in the case of edta, while a chain-terminating reaction such as Reaction (21), or possibly the formation of a ligand radical as described above, becomes more important for edda and nta. Winterbourn [ 1341 also observed more damage to deoxyribose from aqueous iron(I1) and hydrogen peroxide than in a similar experiment with iron(I1) edta, and suggested that iron(1V) was an intermediate during the oxidation of deoxyribose. The fourth reaction is the reduction of hydrogen peroxide by iron(I1) atp or iron(I1) utp in the presence of excess hydrogen peroxide [32]. We presented kinetic evidence that an iron(I1) atp (or utpEhydrogen peroxide intermediate, presumably iron(1V) atp, may yield the hydroxyl radical or react with another hydrogen peroxide to form a bound superoxide. This iron(II1)-superoxide complex subsequently oxidises another iron(I1) atp (or utp) to yield net two iron(II1) atp (or up) per hydrogen peroxide. This mechanism explained our results under conditions of excess hydrogen peroxide. In contrast, Yamazaki and Piette [114,135] argue on the basis of spin-trapping experiments that an intermediate may be formed when the iron(I1) complex is present in excess. This intermediate is observed when edta is the chelating agent, but not with dtpa. They present evidence that such an intermediate is capable of oxidising

18

ethanol, benzoate and tert-butanol, histidine, formate and mannitol. At present we cannot reconcile these different observations. We favour the concept of a higher oxidation state as an intermediate as shown in Scheme 1 (section 3.4): if no scavenger reacts with such an intermediate, the hydroxyl radical will be formed. This is clear from the reaction mechanism published in 1988 [2] and from our reviews [123,127]. A similar view has been expressed in the case of the reaction of copper(1) with hydrogen peroxide [ 1201. Nevertheless, Halliwell and Gutteridge [40,136] argue that either the hydroxyl radical or a higher oxidation state is formed. Since hydroxylation does take place and DMPO-OH adducts are formed, they concluded that iron(1V) compounds are never products of the Fenton reaction [40,136]. Such a simplistic view is incorrect. In the case of the copper(1) ion the situation seems to be more straightforward. The intermediate Cu'H202 was found to react directly with organic scavengers, while in the absence of organic compounds the hydroxyl radical is formed [ 1201. This mechanism explains the small number of cycles observed for the decomposition of hydrogen peroxide by aqueous copper [ 137,1381. The reaction of Cu'( 1, lO-phenh with hydrogen peroxide has been studied in detail [ 1191. The oxidation of copper(1) proceeds with a rate constant of 1.1 x lo3M-' s-' and about equal amounts of HO' and Cu1H2O2(phen)2 are formed. As discussed above, formation of a cupryl intermediate can be ruled out on theoretical grounds.

4. Speciation and effectiveness in promoting oxyradical damage Under physiologically relevant conditions iron is, and copper is likely to be, bound to ligands other than water. In the above sections we have seen that ligands influence rate constants, and that chelating agents that allow access to the metal react faster with hydrogen peroxide. For a catalyst to be effective it should not be inactivated by the products formed. For the Fenton reaction this means that the hydroxyl radical should not react with the ligand, and that the other product, the hydroxide anion, should not be able to displace the multidentate ligand from the iron(II1) ion. Ligands also determine reduction potentials and this can be used to generate conditions to prevent the Fenton reaction. For instance, the reduction potential of the Fe(II1)-/Fe(II)desferrioxamine couple is -0.45V as follows from the difference in stability constants between iron(I1I)- and iron(I1)-desferrioxamine [ 1391. Experimental support for this value exists [ 1401. Ferro-oxamine can reduce hydrogen peroxide and oxygen, but common reductants such as ascorbate and superoxide cannot re-reduce iron(II1) desferrioxamine. Thus, desferrioxamine can be used to prevent the Fenton reaction because iron desferrioxamine cannot act as a

19

catalyst. There is one exception. The methylviologen or paraquat radical is as reducing as ferrooxamine, E0(PQ2'/PQ") = -0.448 V [141], and for that reason desferrioxamine does not protect in systems where this radical occurs [ 1421. Replacement of ligands by desferrioxamine takes several minutes, as shown in a lipid peroxidation study [ 151. One can also add 1,lO-phenanthroline or bipyridine to stabilise the iron(I1) redox state [E0(Fe3+-/Fe2+(phen)3 = 1.14 V, that of the bipyridyl complex is 1.1 1 V] and prevent the reduction of hydrogen peroxide. When initially added to a solution containing iron(III), these ligands may cause some damage because of the oxidising properties of the metal complex. Alternatively, one can use dtpa when these ligands are undesirable. Dtpa allows the reduction of hydrogen peroxide (Table l), but the reduction of iron(II1) dtpa by superoxide is too slow (
20

radical reactions [ 1491 makes this feasible. Furthermore, with the help of the compilations of Martell and Smith [ 139,150-1 541 researchers can design experiments to ensure that enough ligand is present to form a well-defined complex. Scavengers should be tested at more than one concentration. Hopefully, it will be possible in this way to unravel the mechanisms of oxyradical damage.

Acknowledgement The author is grateful to Dr. P.L. Bounds and Dr. D. Bartlett for critical comments during the preparation of this manuscript. Supported by grants from The Council for Tobacco Research - USA, Inc. and by the National Institutes of Health (GM48829).

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