Redox-Active Metals: Iron and Copper

In: Principles of Free Radical Biomedicine. Volume 1 ISBN: 978-1-61209-773-2 Editors: K. Pantopoulos and H. M. Schipper © 2012 Nova Science Publishers, Inc. The exclusive license for this PDF is limited to personal website use only. No part of this digital document may be reproduced, stored in a retrieval system or transmitted commercially in any form or by any means. The publisher has taken reasonable care in the preparation of this digital document, but makes no expressed or implied warranty of any kind and assumes no responsibility for any errors or omissions. No liability is assumed for incidental or consequential damages in connection with or arising out of information contained herein. This digital document is sold with the clear understanding that the publisher is not engaged in rendering legal, medical or any other professional services.

Chapter 5

Redox-Active Metals: Iron and Copper Willem H. Koppenol* and Patricia L. Bounds Institute of Inorganic Chemistry, Department of Chemistry and Applied Biosciences Eidgenössische Technische Hochschule, CH-8093 Zürich, Switzerland

1. Introduction Dioxygen† [1, 2] (O2) is simultaneously necessary for life and a source of potential harm [3, 4]. Harmful reactions, in which partially reduced oxygen species (PROS)‡ such as superoxide (O2 ) and hydrogen peroxide (H2O2) may be formed, are, in part, mediated by redox-active metals, the most physiologically relevant of which are iron and copper. These metals are essential for the transport, storage, and activation of O2, and for electron transfer, e.g., during respiration (see also Vol. II, Chapters 19-20). Along the reductive pathway of O2 to water, partially reduced and potentially reactive intermediates are sequentially made and managed by metalloproteins [5]. Although the redox reactivity of the metals in these proteins *

Email: [email protected] Nomenclature. Formula, systematic name and still allowed trivial name in italics: O 2, dioxygen, oxygen; O2 , dioxide(•1 ) or dioxidanidyl, superoxide; HO2 , hydridodioxygen(•), dioxidanyl, or hydrogen dioxide (hydroperoxyl or perhydroxyl are obsolete); H2O2, dioxidane, hydrogen peroxide; HO , hydridooxygen(•) or oxidanyl, hydroxyl; O , oxide(• ) or oxidanidyl; OCl , oxidochlorate(1 ), hypochlorite; HOCl, hydroxidochlorine, hypochlorous acid; NO , oxidonitrogen(•) or nitrogen monoxide (nitric oxide is obsolete); ONOO , oxidodioxidonitrate(1 ), peroxynitrite; ONOOH, hydroperoxidooxidonitrogen, peroxynitrous acid; Fe3+ or Fe(III), iron(3+) or iron(III) (ferric is obsolete); Fe2+ or Fe(II), iron(2+) or iron(II) (ferrous is obsolete); Cu+ or Cu(I), copper(+) or copper(I) (cuprous is obsolete); Cu2+ or Cu(II), copper(2+) or copper(II) (cupric is obsolete); FeO2+, oxidoiron(2+) or oxidoiron(IV) (ferryl is obsolete)[1,2]. The abbreviations edta, dtpa and nta, and atp refer to the metal chelators ethylenediaminetetraacetate, diethylenetriaminepentaacetate, nitrilotriacetate and adenosinetriposphate, respectively; these abbreviations are not capitalised.[2] NTBI is non-transferrinbound iron; cp20 is 3-hydroxy-1,2-dimethylpyridin-4(1H)-one, and icl670, 4-[3,5-bis(2hydroxyphenyl)-1H-1,2,4-triazol-1-yl]benzoic acid. ‡ We use the acronym PROS rather than ROS (for “reactive oxygen species”); PROS is not generally meant to include O2, which, strictly speaking, is also a reactive species. Analogously, we apply the acronym PONS for “partially oxidized nitrogen species” instead of RNS for “reactive nitrogen species”. †

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Willem H. Koppenol and Patricia L. Bounds

is harnessed and directed by the protein and/or co-factor to which it is bound, there can be some non-productive release of O2 or H2O2. One approach that nature uses to circumvent non-productive reactions of oxygen is to employ more than one redox-active metal ion, sometimes stabilized by one or more non-redox-active metal ions, for the transfer of multiple electrons over a short time frame, e.g. a copper and an iron ion for the reduction of O2 by cytochrome c oxidase [6], and four manganese ions for the oxidation of water in photosystem II [7]. Even so, oxyhemoglobin [8, 9] and oxymyoglobin [10] reportedly release O2 , and xanthine oxidase [11], and cytochromes P450 [12] are known to produce H2O2. Another source of O2 in vivo is the reduction of O2 by electrons that leak from the mitochondrial electron transport chain; in fact, 1 – 4% of all O2 consumed during respiration may be reduced to O2 in this fashion [13, 14] (Vol. II, Chapter 15). It is now well established that PROSs, including peroxynitrite (ONOO ), play roles in a stunning array of diseases [15, 16]. One function of the protein environment of a metalloprotein is to modulate the reactivity of the metal ion; the potential for harmful reactivity likely to be catalyzed by free metal ions such as iron and copper outside of the protein milieu is the topic of this review. In vivo, excess free metal ions are sequestered in specialized storage proteins ― iron in ferritin [17] (Vol. II, Chapter 19) and copper possibly in metallothioneins (Vol. II, Chapter 20), while circulating iron is captured by transferrin [18] and copper by ceruloplasmin [19]. However, transfer of metals from storage to functional metalloproteins and the turnover of these proteins likely gives rise to low concentrations of low-molecular weight metal ion complexes [20, 21]; this pool of iron and copper may increase as a consequence of oxidative stress.

2. Redox Active Iron and Copper In Vivo The role played by iron and copper in diseases of ageing and oxidative stress has been recently reviewed [22]. Both iron and copper ions “redox cycle”, i.e. shuttle between two oxidation states via transfer of a single electron (Figure 1), as part of their normal functions in vivo. Redox cycling is also widely believed to be involved in the generation of damaging PROS; “free” iron and copper, actually free in solution or non-specifically bound to albumin or lowmolecular-weight ligands, are invoked as being particularly damaging. However, the link between pools of free transition metals, PROS, and disease is tenuous, and the evidence in the literature is largely circumstantial. Further, the speciation, that is the ligands, of “free” iron and copper ions involved in disease states have not been precisely identified. Diseases in which excess iron plays a role include hemochromatosis and iron-loading anemias, such as thalassemia, sickle cell disease and myelodysplasia [23] (Vol. III, Chapter 2). When normal mechanisms of iron intake and elimination become imbalanced, the binding capacity of ferritin and transferrin are exceeded, and pools of free, or non-transferrin-bound, iron (NTBI) increase (Vol. II, Chapter 19). The NTBI, which has been detected in the serum of hemochromatosis [24], sickle cell anemia and thalassemia patients [25], may consist in part of iron(III)-citrate complexes [26]. Although citrate has been identified as a ligand in NTBI, the exact composition of the complex or complexes existing in vivo remain elusive [27]. Citrate is present in the serum at a concentration of 0.1 mM, and the clinical range of NTBI concentrations varies between 0 to 10 M [28]. The aqueous speciation of iron(III) citrate is highly dependent on the iron/citrate molar ratio and, under the clinical conditions described,

Redox-Active Metals: Iron and Copper

93

several complexes have been shown to be possible [29]. The mononuclear Fe(cit)25 complex predominates over the clinically observed iron/citrate molar ratios, and, additionally, the multinuclear complexes Fe3(cit)33 and Fe3(cit)47 become relevant at higher NTBI concentrations [30]. Thus, the term NTBI may encompass several different species, of which one or more could be redox-active. Altered concentrations of copper, zinc and iron ions in the brain are a feature of Alzheimer‟s disease, and copper is widely thought to contribute to the oxidative stress model of Alzheimer pathophysiology [31] (see also Vol. III, Chapter 10). Copper has been shown to bind tightly to amyloid- peptide (K = 10–11 M), and the copper-amyloid- complex oxidizes cholesterol to cytotoxic products [32, 33]. Copper also plays a role in Parkinson‟s disease, atherosclerosis, diabetes, and other age-related diseases, as reviewed by Brewer [22].

3. Biological Oxidants Reactive species formed in vivo that are thought to be physiologically relevant include O2 , nitrogen monoxide (NO ), the hydroxyl radical (HO ), the trioxocarbonyl radical (CO3 ), nitrogen dioxide (NO2 ), H2O2, ONOO , and hypochlorite (OCl ) (see also Chapters 2-4). The first four species in the list each possess an unpaired electron and are, therefore, called radicals; H2O2, ONOO and OCl have no unpaired electrons and are, therefore, not radicals§. It is important to recognise that radical species are also functionally produced in vivo, e.g., the nitric oxide synthases generate NO , a diatomic radical species that plays vital roles in neuronal cell communication, endothelial vasodilation, and immune response [34, 35]. Of all the biological oxidants listed above, HO is the most destructive: HO is a promiscuous oxidant that reacts with small organic molecules, proteins, nucleic acids, and lipids by abstraction of hydrogen or addition to a double bond to form a radical species that can undergo further transformations. The HO has a high positive electrode potential|| (E° (HO ,H+/H2O) = +2.31 V at pH 7) [36], and the second-order rate constants (kobs = 109 – 1010 M 1s 1) [37] are in a range that indicates diffusion-controlled reactions. In contrast, H2O2 does not react directly with most organic compounds, and any toxicity attributed to H2O2 §

||

The term “radical”, which was used in the past to refer to a group that is part of an organic molecule, e.g., “methyl radical,” has been replaced with substituent; in former usage, the term “free radical” referred to an unattached group, usually containing an unpaired electron. In current usage, the word “radical” indicates that the group possesses an unpaired electron, and the word “free” is no longer included. The values given are relative to the normal hydrogen electrode, a platinum electrode in a solution that is 1 molal in H+ (1 gram H+ per kg of water, essentially the same as 1 molar) and in equilibrium with an atmosphere of 100 kPa of H2, that, by definition, has a value of 0 V. If we use O 2 and O2 as an example, the electrode potential of the couple O2/O2 refers to the reaction: O2 + e O2 , with a value of 0.35 V, relative to the normal hydrogen electrode. The electrode potential was formerly called a reduction potential. This value would be measured if one could connect a stable solution of O 2 (1 molal) in equilibrium with an atmosphere of O2 (100 kPa) with the normal hydrogen electrode by way of two platinum electrodes connected to a volt meter, and a salt bridge between the solutions to allow ions to flow between the solutions. Electrons would flow from the negatively charged electrode in the O2/O2 solution – producing O2 – to the positive electrode in the H+/H2 solution where H2 would be formed. It should be clear, though, that a stable solution of O2 does not exist; instead, the electrode potential is measured by indirect means. An oxidation potential refers to the reaction O2 O2 + e and the sign is reversed: +0.35 V.

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would depend on the presence of redox-active transition metal ions, iron and copper, the topic of this Chapter. It is known that NO , O2 , and H2O2 are produced enzymatically, and the remaining species listed are formed from these (Figure 2). In phagocytic cells, i.e., neutrophils and macrophages, O2 generated by NADPH oxidase reacts to produce H2O2 (Reaction 1) and ONOO (Reaction 2) [38]. Also in neutrophils, H2O2 is a substrate in the reaction of myeloperoxidase to generate OCl (Reaction 3). The reactive ONOO and OCl synthesized in phagocytic cells are part of the host organism‟s defence arsenal to combat pathogenic organisms [39] and destroy foreign objects [40]. Reaction 4, the reaction of ONOO with carbon dioxide (CO2), yields mostly NO3 and CO2, but also the oxidizing radicals CO3 and NO2 [41, 42]. 2 O2 + 2H+

O2 + H2O2

(1)

O2 + NO

ONOO

(2)

H2O2 + Cl

OCl + H2O

(3)

ONOO + CO2

CO3 + NO2

(4)

Hasc–

asc – + H+

NTBI-Fe3+

NTBI-Fe2+

H2O2 + H+

H2O + HO Figure 1. Redox cycling by iron.

+

L-Arginine + 1.5 (NADPH + H ) + 2 O2

PROS

+

L-Citrulline + 1.5 (NADP ) + 2 H2O

NOS

CO2

NO

CO3 – + NO2

ONOOCO2–

ONOO–

CO2 + NO3–

NADPH + e– NADP– + H+

O2

NOX

O2

–

HO2 H

Figure 2. Generation and interconnectedness of PROS.

H2O2

SOD

+

O2

–

O2

OCl–

MPO

Cl–

H2 O


95

The electrode potential of O2 , E°(O2/O2 ) = –(0.18 + 0.02) V, or –(0.35 + 0.02) V (pO2 = 0.100 Mpa) indicates that O2 is only mildly reducing; O2 is, indeed, not very damaging [43]. Reaction 1, the disproportionation of O2 is catalyzed by superoxide dismutases [44] (SODs; see Vol. II, Chapter 5). At pH 7, O2 is 0.5% protonated (HO2 ), and the protonated form also acts as an oxidant: E° (HO2 ,H+/H2O2) = 1.04 V [45]. The electrode potentials indicate that the oxidation of O2 – by HO2 is thermodynamically likely, and the reaction indeed proceeds rapidly (k = 1.0 • 108 M 1s 1) [46] even in the absence of SOD. The importance of NO as a biological signaling molecule [47] has emerged in the past two decades. In our view, NO itself is not damaging per se aside from its affinity for iron(II) which may lead to inhibition of hemoproteins with an open coordination site. The very rapid reactions of NO with oxymyoglobin and oxyhemoglobin (k = 4.4 • 107 M 1s 1 and 8.9 • 107 M 1s 1, respectively, relative to heme concentration) produce the corresponding metmyoglobin and methemoglobin (and nitrate) [48] in which the heme iron is oxidized and cannot bind O2. In theory, NO2 can be formed via Reaction 5, the autoxidation of NO ; this reaction, however, requires two NO and one O2 [49-53], and the mechanism is complex [54]. Reaction 5 is likely to be slow (hours to days) in vivo, given the estimated physiological concentrations of NO and O2 [55]. In contrast, NO diffuses from tissue to the nearest red blood cell, where it is scavenged by hemoglobin more rapidly than it can react with O2 [56]. Thus, the lifetime of NO in vivo is determined by reactions with O2 and hemoglobin, and Reaction 5 can be neglected. 2 NO + O2

2NO2

(5)

The reaction of NO with O2 to form ONOO– (Reaction 2) is diffusion-controlled (k = 1.6 • 1010 M 1s 1) [57], and ONOO– is recognized as an important biological oxidant. At physiological pH, ONOO– is largely protonated (pKa = 6.8) [58]; ONOOH reacts much more slowly and selectively than HO [59] with biological targets. Evidence for formation of ONOOH in vivo is largely based on immunodetection of nitrotyrosine residues [60]; it should, however be noted that nitrotyrosine may also be formed by peroxidase-catalyzed reactions of nitrite and H2O2. Tyrosine nitration blocks phosphorylation and, thus, impacts on the signaling functions of tyrosine residues [61]. Furthermore, nitrotyrosine can, at least in vitro, accept an electron from ascorbate and then reduce O2 [62], thereby contributing to the formation of O2 –. The main reaction of ONOOH is isomerisation to NO3–, but it is widely believed that ONOOH undergoes homolysis to at least some extent to yield NO2 and HO [63-65]. Estimates of the yield of homolysis range from as low as ca. 1% [66] to as high as ca. 40% [67]. We have sought but failed to find evidence to support a mechanism in which homolysis accounts for more than a very low percentage of the decay of ONOOH [68-72]. The radicals produced in the reaction of ONOO– with CO2 [73], NO2 and CO3 [41, 42] are strongly oxidizing species: (E° (NO2 /NO2 ) = +1.04 V) [74] and (E° (CO3 /CO32 ) = +1.58 V) [75, 76]. These radicals have been implicated in ONOO–-mediated nucleic acid oxidation [77] and tyrosine nitration [78]. Finally, the OCl– produced by myeloperoxidase in activated phagocytic cells reacts with many biomolecules, including nucleotides, DNA and low-molecular-weight thiols [79], as well as protein thiol [80] and amino residues [81].

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4. Redox Chemistry and Biochemistry of Iron: The Historical Context The involvement of transition metal ions in oxidative stress pathologies is often attributed to “Fenton chemistry”, and the Fenton reaction, reduction of H2O2 by iron(II) to generate HO (Reaction 6), has long been of interest to the free radical community. H. J. H. Fenton discovered the reaction that bears his name serendipitously [82]. While still an undergraduate student, Fenton was shown a violet coloured solution that had been obtained by a fellow student who was mixing reagents, including tartaric acid, at random, and later reproduced the reaction and published it as a test for tartaric acid in 1876 [83]. A full account, including the structure of the oxidized product, 2,3-dihydroxymaleic acid, was published in 1896 [84]. Although Fenton used iron(II) and H2O2 to modify organic compounds, he was not aware that HO• is formed in the reaction. Two years after Fenton‟s death, Fritz Haber and Richard M. Willstätter [85] published a paper on radical chain reactions in organic chemistry and biochemistry, in which they attributed the action of catalase on H2O2 to the initiation of a radical reaction (Reaction 6, the Fenton reaction) [86]. Later, Haber and his assistant Joseph Weiss [87, 88] investigated the decay of H2O2 by iron salts at low pH and concluded that, since more than 1 H2O2 per 2 Fe2+ is consumed when H2O2 is present in excess, the mechanism involves a radical chain reaction (Reactions 7 and 8), with chain termination by Reaction 9. Fe2+ + H2O2

Fe3+ + HO + HO•

(6)

HO• + H2O2

H2O + O2• + H+

(7)

O2• + H+ + H2O2 Fe2+ + HO• + H+

O2 + HO• + H2O Fe3+ + H2O

(8) (9)

By the late 1940s, the mechanism proposed by Haber and Weiss for the iron-catalyzed decomposition of H2O2 had been criticised by Philip George, who showed that O2•– does not react with H2O2 [89]. An alternative mechanism at low pH proposed by George and co-workers in 1949 [90] on the basis of product analyses is summarised in Table 1. Note that O2 occurs in these reactions as HO2 (pKa = 4.8) [43]. More extensive proposals for the mechanism were described in two publications in 1951 [91, 92], the first of which contains a reference to Fenton's first full report of the oxidation of tartaric acid [93]. The rate constants were determined later. The mechanism published by George and coworkers was corroborated in 1985 [94]. It should be noted that all of these investigations of iron-catalyzed decomposition of H2O2 were carried out at low pH and are generally not physiologically significant. A new field of study, free radical biochemistry, was created in the wake of the discovery in 1969 of an enzymatic function for hemocuprein [95], namely catalysis of the disproportionation of O2 (Reaction 1) [96]. The rate constant for the copper/zinc SOD (CuZnSOD) reaction, which is governed by electrostatic guidance of O2 to the active site [97, 98], indicates that the reaction is close to diffusion-controlled (kcat 1 2 109 M 1s 1) [99, 100]. The majority of

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CuZnSOD, ca. 5–10 M, is found in the cytosol of the cell; some is also found in the mitochondrial intermembrane space [101]. In 1973, Fe- and Mn-containing SODs were discovered [102, 103]: the mitochondrial matrix contains manganese SOD (MnSOD), and bacteria contain a structurally related Fe-enzyme. A tetrameric CuZnSOD manages O2 in the extracellular space [104]. Given that Reaction 1 uncatalyzed is rapid (k = ca. 106 M 1s 1 at neutral pH), and that SODs are widespread, the conclusion is warranted that O2 is reactive and, therefore, must be dangerous. Unfortunately, George‟s work from the 1940s was not consulted, and, in the 1970s, reduction of H2O2 by O2 (Reaction 8) from the Haber-Wilstätter mechanism was proposed [105] as a possible harmful reaction from which the organism needed to be protected. At this point in history, Reaction 8, in which the relatively innocuous O2 is converted to the far more reactive HO radical, became known as the Haber-Weiss reaction#. Again, it had to be pointed out in other studies that Reaction 8 is too slow to be significant [106-111]. When it was realized that O2 does not reduce H2O2, iron complexes were invoked to act as catalysts, as in the mechanism of George et al., and Reactions 6 and 9 became known as the “Fenton-catalyzed Haber-Weiss reaction.” Later, it was realised and accepted that other reductants, e.g., monohydrogen ascorbate, are more likely candidates to reduce Fe3+ complexes in vivo.

5. Biological Relevance of the Fenton Reaction The toxicity of iron, and, by implication, copper, originates from its ability to reduce peroxides, via Fenton chemistry. Although George and coworkers provided a mechanism for the decay of H2O2 catalyzed by iron at low pH, it is not clear that the mechanism has any meaning at neutral pH where aqueous Fe3+ ions are not available for reaction because they have formed hydroxide, phosphate, or other complexes. If such Fe3+ complexes are present in solution, are they reduced by O2 , or monohydrogen ascorbate, and are the corresponding Fe2+ complexes likely to be oxidized by H2O2 or organic peroxides? Table 1. Rate constants for the reactions in the mechanism of iron-catalyzed decomposition of H2O2 at low pH [89] Reactions Fe2+ + H2O2 + H+ Fe3+ + H2O + HO• HO• + H2O2 H2O + HO2• Fe2+ + HO• + H+ Fe3+ + H2O 2+ + Fe + HO2 + H Fe3+ + H2O2 Fe3+ + HO2 Fe2+ + O2 + H+

#

Reaction No. 9 7 10 11 12

k (M 1s 1) 41.5 2.7 • 107 4.3 • 108 1.2 • 106 2.0 • 104 (pH 1)

It is not clear why Reactions 7 and 8 have become known as the Haber-Weiss cycle rather than the HaberWillstätter cycle; it may be because the often-quoted paper by Haber and Weiss is in English and, thus, more accessible to the international research community, while the language of the original paper of Haber and Willstätter is German. Haber and Willstätter 87 is cited by Haber and Weiss 88 but neither paper refers to Fenton 84, although Reaction 6 is the Fenton reaction.

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The thermodynamic requirements for a metal complex in aqueous solution at neutral pH to redox-cycle are simple: when O2 – is the reductant, the metal-chelate must have a electrode potential between E°‟(O2/O2 –), –0.35 V and E°‟(H2O2/ OH,H2O), +0.39 V [5, 112]. The window of redox-opportunity (0.74 V) is therefore wide. For both iron and copper, redox cycling involves transfer of a single electron; thus, only one-electron electrode potentials may be considered. Potentials for two-electron oxidations, e.g., two-electron electrode potential of the NAD+, H+/NADH couple (–0.32 V) [113, 114], should never be considered in discussions of redox-cycling. Such a reaction is kinetically not feasible, as it would require the simultaneous collision of three reactants, NADH and two iron complexes, which is unlikely under the dilute concentrations found physiologically. Although O2 – has been shown to support hydroxylation of benzoic acid catalyzed by iron chelates [115], these studies were carried out with chelating agents that, with the exception of citrate [116], are not physiologically relevant. It is also important to recognize that, in most tissue compartments, the steady-state concentration of O2 – is maintained at a vanishingly low level by SODs [117]. The rate constant for reduction of iron(III)-edta† by O2 –, determined directly by pulse radiolysis, is ca. 1 • 106 M 1s 1 at physiological pH [118], three orders of magnitude smaller than the catalytic rate constant of CuZnSOD; rate constants for the reaction of iron(III)-dtpa† and -atp were much less than 1 • 106 M 1s 1 and could not be determined. Not surprisingly, it is now felt that monohydrogen ascorbate is a more likely physiological reductant of iron(III) complexes. Given that the electrode potential of the asc /Hasc couple is +0.28 V [119], the redox window for redox cycling involving ascorbate is much smaller, i.e., 0.11 V. However, for a typical serum concentration of 50 M monohydrogen ascorbate, the concentration of the ascorbyl radical is likely to be in the nanomolar range. Thus, the electrode potential for the ascorbyl/monohydrogen ascorbate couple is, according to the Nernst equation, ca. 0.18 V less than +0.28 V, which would widen the redox window to 0.29 V. Similar considerations of the H2O2,H+/HO , H2O couple increases the electrode potential from +0.39 V to ca. +0.9 V, an adjustment that results in a much larger window of 0.8 V, from +0.1 V to +0.9 V [120]. A number of studies with iron complexed to citrate, atp or aminopolycarbonates, such as edta, dtpa and nta†, suggest that the Fenton reaction, and its equivalent with organic peroxides, is thermodynamically feasible [121] and does proceed. Rate constants of the order of 102 to 105 M 1s 1 have been reported for the Fenton reactions of these various iron(II) complexes [122131]; there is, however, no good agreement between the published rate constants for a given complex. The reaction, which requires binding of H2O2 to iron(II), generally proceeds faster when more exchangeable water molecules are bound to iron(II) [132], i.e., when more ligand sites on the iron are available. However, all these reactions are much slower than that of H2O2 with catalase, and, thus, not likely to be physiologically significant.**

**

It is not sufficient to simply compare rate constants; the relative concentrations of the Fe2+ complex and catalase must be considered as well. It is the product of the concentration of the iron complexes or catalase with the respective rate constants for the reaction with H2O2 that must be compared. Catalase is present at ca. 2 M, and the concentration of redox-active iron is, at most, a few M. In the following, the concentration of the Fe2+ complex 5 M is used as an example to demonstrate that the Fenton reaction with Fe2+ complexes is outcompeted by catalase. This example is, of course, valid only for homogeneous solution: kFenton[Fe2+] = (1 • 104 M 1s 1)(5 • 10 6 M) = 0.05 s 1


99

In vitro studies performed under physiological conditions indicate that HO• is not the only oxidant formed during the Fenton reaction [121, 133-135]. Depending on the conditions, the oxidant may be (1) oxidoiron(IV), a higher oxidation state of iron, as suggested during the 1930s (Reaction 13) [136], (2) H2O2 bound to iron(III), or (3) HO•. The pattern of hydroxylation of salicylate by iron(II)-edta and H2O2 compared to that by HO• generated by radiolysis of water suggests that HO• formation is likely [137]. Isotope/ESR studies established that the oxygen in either the HO• or oxidoiron(IV) is from H2O2 [138]. Fe2+ + H2O2

FeO2+ + H2O

(13)

What is the likelihood that the Fenton reaction occurs in vivo? Assuming that monohydrogen ascorbate is the source of electrons in vivo, there are number of problems with the hypothesis that iron is a catalyst in oxidative processes: (1) We do not know the composition/structure of the iron complex (or complexes) in vivo reduced by monohydrogen ascorbate and subsequently oxidized during the Fenton reaction, nor do we know any rate constants for the reactions. (2) The rate constants for the Fenton reaction (102 105 M 1s 1) that have been determined are remarkably unimpressive when compared to that for the reaction of catalase with H2O2 (3.5 107 M 1s 1) [139]. Additionally, H2O2 is scavenged by glutathione peroxidase [140] and peroxiredoxin [141]. (3) The Fenton reaction is even slower in the presence the NO , which is present in vivo and which binds to iron(II) [142]. (4) The hypothesis does not include a protective role for SOD, but only for proteins that remove H2O2, which is contrary to experimental observations that SOD alone can protect tissues from oxyradical damage. Given that most biomolecules are unreactive toward O2 –, concentrations of ca. 5–10 M CuZnSOD in the cytosol would seem higher than necessary to protect against damage from O2 –. Clearly, SOD does protect, but not from Fenton chemistry. Beckman et al. [143] wrote in 1990 that “generation of strong oxidants by the iron-catalyzed Haber-Weiss reaction is not an entirely satisfactory explanation for SOD-inhibitable injury in vivo” and proposed that NO reacts with O2 – to form ONOO–, and that formation of ONOOH is kinetically far more feasible than the “iron-catalyzed Haber-Weiss reaction.” Earlier in this Chapter, we described ONOO as an important and selective oxidizing and nitrating species. Thus, SOD protects by eliminating O2 – and prevents, thereby, formation of ONOO . The diffusion-controlled rate constant for the reaction of O2 – with NO [57] necessitates the micromolar concentration of SOD observed in vivo. Where, then, may the Fenton reaction play a role? In the acidic environment of the lysosome, where ferritin is decomposed, the pH is between 4 and 5; here, the concentration of chelatable iron is relatively high (ca. 15 M), catalase is absent, the coordination of iron is altered and the iron may be reduced due to the presence of cysteine. Under conditions of oxidative stress, lysosomes may be damaged and ruptured, which would release relatively high concentrations of redox-active iron [144, 145] into the cytosol (Vol. II, Chapter 16). Ferritins released into the bloodstream during inflammation and autoimmune diseases may contribute to locally high iron concentrations that lead to apoptosis [146]. Fenton chemistry-derived oxidative damage can be prevented by chelation of the iron with desferrioxamine (dfo) [146-148].

kcatalase[catalase]

= (3.5 • 107 M 1s 1)(2 • 10 6 M) = 70 s

1

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The extracellular NTBI levels may contribute to intracellular labile iron pool (LIP) levels (Vol. II, Chapter 19) and play a role in oxyradical damage [149], or be a symptom of iron overload. Any form of mobile iron poses a threat to the liver and the heart, where excess iron accumulates, and iron chelation therapy is employed to promote excretion. The „wish list‟ for an effective iron chelator has been described as: „stable in vivo, non-toxic, rapidly excreted, preferably orally-active, and produced with a low synthetic cost‟ [150]. Toxicity is ascribed to Fenton chemistry, thus, the chelate ought not be redox-active: the three chelating agents used clinically, dfo, cp20 and icl670, bind iron(III) far more tightly than iron(II), which results in very negative standard electrode potentials ( 0.45 to 0.60 V [120, 151-153]) well outside the window of redox opportunity, such that monohydrogen ascorbate cannot reduce these iron(III) complexes. The low electrode potentials also imply that dfo, cp20 and icl670 bind iron(II) at micromolar concentrations only in part or not at all, which has been experimentally verified for cp20 [154]: when iron(II) is oxidized, it is complexed by cp20, thereby preventing reduction and redox cycling. Although copper is less ubiquitous than iron in living systems, copper and iron proteins have similar functions, including electron transfer and O2 binding and activation. In humans, copper is used for electron transfer reactions in the enzymes CuZnSOD and cytochrome c oxidase. During the 1980s, it was recognized and demonstrated that free copper ions can, in theory, redox cycle and promote reactions analogous to those catalyzed by free iron [155158]. In in vitro studies under acidic conditions, copper(I) was shown to form an intermediate Cu+H2O2, which reacts with organic scavengers, however, not by a Fenton mechanism [159]. Other investigators [160] concluded on the basis of pH and substrate concentration dependence experiments that no HO is formed during the reaction of copper(I) with H2O2. Studies with red blood cells show that monohydrogen ascorbate can reduce copper complexes to initiate reactions similar to those described for iron [161]. However, the concentration of free copper ions in the cell is extremely low, only ca. 1 copper ion per cell [162], and the possibility that copper participates in oxyradical damage reactions in vivo is vanishingly small. The disproportionation of O2 – catalyzed by free copper ions is faster by a factor of 4 than that catalyzed by SOD [163, 164]; the inclusion of edta to bind free copper(II) prevents the reaction with O2 –. It is important to note that edta, nta and dtpa, common metal chelation agents that are often used as models for physiological low-molecular-weight iron complexes in biochemical studies, may enhance the Fenton activity of iron at neutral pH. Depending on the complex, one or more coordination sites of iron are occupied by water, which may give way to H2O2 [132]. In the case of copper, the complexes formed with these ligands are essentially inert at physiological pH, and redox chemistry is prevented. Thus, addition of chelators can produce differential effects on redox behavior depending on the target metal, e.g., edta and nta enhance iron-mediated damage to chromatin but protect against copper-mediated damage [165]. Similar differential effects of edta on oxidative reactivity in cerebrospinal fluid and brain tissue was recently reported [166], and these authors correctly point out that different chelation strategies are required for sequestration of iron – “the main characteristics of an iron sequestrant should be its ability to compete with metal ligands naturally present in CSF” – and copper – “copper requires application of a strong and efficient chelator”.


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6. Cautionary Remarks Given the low micromolar concentrations of iron and H2O2 found under normal physiological conditions and the rate constants for Fenton reactions with iron bound to physiologically relevant ligands, it is kinetically unlikely that iron causes damage, and the same may hold for copper. Iron and copper may contribute to oxidative processes under conditions of higher local concentrations of these ions. The destruction of ferritin (and by extension ceruloplasmin) in lysosomes may provide such concentrations. NTBI in blood is an indicator of excess iron; this form of iron does not appear to be very redox-active, but can be considered dangerous since it is a form of mobile iron that may be deposited in the liver or in the heart. To determine the rate constant for a reaction of physiological interest, it is usually necessary to use reagents at concentrations that are significantly higher than physiological conditions; the rate constant thus determined can be extrapolated to the dilute conditions found in vivo to estimate how fast the process is in reality. However, a rate constant determined in vitro is valid for homogeneous solution, whereas the many reactions going on inside the cell take place under more heterogeneous conditions. Thus, the observation that common HO∙ scavengers fail to protect against oxidative damage in vivo has been attributed to formation of HO∙ at locations within cellular components that are poorly accessible to the scavenger. Alternatively, formation of oxidoiron(IV), which might react less rapidly than HO∙ with scavengers, has been postulated. Results obtained from experiments performed with cells to which relatively high concentrations of metal ions and/or H2O2 have been added cannot be realistically extended to prove that oxidative damage occurs under normal physiological conditions, where the antioxidant defence system consisting of monohydrogen ascorbate, glutathione, vitamin E, SODs, catalase, glutathione peroxidase, peroxiredoxin, etc. is intact. Some in vitro experiments can be additionally criticized for having used simple iron salt solutions: e.g., iron(III) salt solutions are very acidic, and, at neutral pH, colloids of iron(III) hydroxide/phosphate and iron(III)oxide, which may simply be toxic rather than catalysts of redox reactions, are formed. Iron(II) autoxidizes rapidly at pH 7 to generate O2∙– and H2O2, which are likely to complicate the interpretation of findings. In short, the speciation of a metal ion determines its thermodynamic and kinetic properties and, by extension, its reactivity. One obstacle to progress in the field of redox metal toxicity is the inability to follow oxidative stress in cells in a time-resolved fashion. Although widely used fluorescence techniques have been touted to do so, the addition of the fluorescent indicator itself may lead to the release of PROS [167, 168].

Conclusions Are redox-active metal ions causative agents in disease, or do diseases cause redox-active metal ions to accumulate? Although metals, especially iron and copper, are clearly implicated in a variety of oxidative-stress-related pathologies, it has not yet been definitively demonstrated whether low-molecular-weight metal ion complexes act to initiate damage or

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are simply liberated as a sequela of oxyradical damage, and it is possible that metals, once released as a result of an oxidative insult, act in a cascade fashion to propagate injury. It is the findings from in vivo studies that most clearly support a role for involvement of redox-active metals in oxidative stress-related pathologies; results from in vitro mechanistic studies based on kinetics have sometimes called those findings into question. Both approaches to the study of free radical biology and medicine have merit and contribute to the ultimate goal of elucidating mechanisms that make sense at the molecular, cellular and organismic levels alike.

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