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Chem. Res. Toxicol. 1996, 9, 828-835
Reactions of Nitric Oxide with Metalloproteins Rafael Radi* Departamento de Bioquı´mica, Facultad de Medicina, Universidad de la Repu´ blica, Montevideo, Uruguay Received October 16, 1995
I. Introduction The interactions of nitric oxide1 (NO) with metalloproteins constitute relevant molecular mechanisms accounting for signal transduction, cytotoxicity, and modulation of free radical chemistry. NO can readily interact with different metal centers in proteins including heme-iron, iron-sulfur clusters, zinc-sulfur clusters, and copper. The reactions of nitric oxide with metal centers in proteins typically result in either the formation of stable nitrosyl-metal complexes and/or redox chemistry. In both types of interaction, NO usually leads to some extent of modification of the protein function. In addition to the analysis of the direct reactions of nitric oxide with metalloproteins, another important point has to be taken into consideration when analyzing the role of nitric oxide in alterations of metalloprotein structure and/or function involved in cytotoxic events. This point is that not always NO directly but in some cases nitric oxide-derived species that are formed in aerobic biological environments are the ultimate reactive species interacting with the metalloprotein. In particular, the formation of peroxynitrite anion (ONOO-) secondary to the almost diffusion-controlled reaction between superoxide and nitric oxide can account for part of the cytotoxic events. In this review, basic information regarding the physicochemical characteristics that make NO a good ligand or reactant for metalloproteins, the interactions between NO and iron-, copper-, and zinc-containing metalloproteins, the differential reactivity of nitric oxide and peroxynitrite toward metalloproteins, and the chemical mechanisms by which nitric oxide can modulate metalloprotein-dependent free radical processes will be provided.
II. NO Interactions with Iron-Containing Proteins a. Electronic Configuration of NO. NO is a free radical species as it contains an odd number of electrons (fifteen) of which seven are located in the outer shell. The molecular orbital electronic configuration of nitric oxide is (Be)2(2pσ)2(2pπ)4(2 pπ*)1 (1). This results in the presence of the last electron in a π* antibonding molecular orbital and an overall bond order between the nitrogen and oxygen atoms of 2.5. NO can interact with * Address correspondence to: Dr. Rafael Radi, Departamento de Bioquı´mica, Facultad de Medicina, Universidad de la Repu´blica, Avda. Gral. Flores 2125, 18000 Montevideo, Uruguay. Fax: (5982) 949563, E-mail:
[email protected]. 1Most of the nitrogen oxides have common names used in the biochemical literature that do not follow the current IUPAC rules for systematic names of inorganic compounds. The IUPAC recommended nomenclature for nitric oxide (‚NO) is nitrogen monoxide, for peroxynitrite anion (ONOO-) it is oxoperoxonitrate(1-), for nitrosonium (NO+) it is nitrosyl cation, for nitronium (NO2+) it is nitryl cation, and for nitroxyl anion (NO-) it is oxonitrate(1-). We will use common names throughout the text due to their widespread use in the biochemical literature.
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both ferrous and ferric iron, but it is a better ligand for ferrous iron since it has an additional electron in the d-orbital (d6) compared to ferric iron (d5), which increases the interactions between iron d-orbitals and the antibonding 2 pπ* electron in NO. By this mechanism, a nitrogen-metal σ-bond is formed by using an unshared electron pair in the nitrogen atom, and a significant degree of π-bond is formed by backbonding between metal d-electrons and the pπ* orbitals of NO (2). Of the total four electrons needed to fulfill the interactions (Fe)dfπ* (NO), three electrons will be provided by the iron and one by nitric oxide. In these iron-nitrosyl complexes nitric oxide usually forms linear complexes with iron, where nitric oxide participates as a three-electron donor, with two electrons participating in the formation of the initial σ-bond of the nitrogen-metal interactions, and the third electron can reduce ferric to ferrous iron leaving NO with a nitrosonium (NO+)-like character. Although association rates of nitric oxide with ferrous iron are typically faster than those for ferric iron, NO can still be a good ligand for ferric iron due to its polarity. In fact, NO is more polar than molecular oxygen and can diffuse better through proteins, such as myoglobin (3). Another important aspect of nitric oxide reactivity with metalloproteins involves its redox properties. NO is neither a strong oxidant nor a strong reductant. The oneelectron reduction of NO to nitroxyl anion (NO-) occurs with a standard one-electron redox potential of +0.38 V while one-electron oxidation to nitrosonium (NO+) is +1.21 V (4). NO usually participates in redox chemistry when interacting with metalloproteins containing iron or copper, but the characteristics of the reaction depend on the redox properties of the metal bound to the protein. It is frequent that the interactions of NO with metalloproteins implicate electron donation. b. General Properties of NO Interactions with Heme Proteins. Long before the discovery of NO as a biologically-produced molecule, it was known that heme proteins could readily interact with nitric oxide. Indeed, early experiments utilized nitric oxide as a spin-ligand probe to study spectral and paramagnetic properties of heme proteins, in particular those of hemoglobin and myoglobin (5-9). The reactions of nitric oxide with different heme proteins share some general characteristics. NO can interact with heme proteins in either the ferrous (Fe2+) or the ferric form (Fe3+). Moreover, it has been recently shown that NO can even react with higher oxidation states of the heme iron such as the ferryl form (Fe4+dO). Kinetic studies have shown that NO reacts with the ferrous state of heme proteins typically faster (107 M-1 s-1) than with the ferric state (102-107 M-1 s-1) (10). Once formed, the ferrous or ferric, heme-nitrosyl complex slowly dissociates, with the dissociation usually faster for the ferric-heme nitrosyl complex. Considering the following equilibrium for either ferrous or ferric heme protein (Hp) interaction with NO: © 1996 American Chemical Society
Forum on Nitric Oxide: Chemical Events in Toxicity kf
Hp + NO h Hp-NO kd
Keq )
[Hp-NO] kf ) [Hp][NO] kd
Chem. Res. Toxicol., Vol. 9, No. 5, 1996 829
(1)
(2)
it becomes apparent that as a consequence of a larger second order rate constant (kf) for combination between heme2+ and NO and a smaller first order rate constant (kd) for dissociation of the heme2+-NO complex, the equilibrium constants (Keq) for the reaction of NO with ferrous heme proteins are larger than those for ferric heme proteins. Heme proteins with an unoccupied 6th coordination position (i.e., hemoglobin, myoglobin, catalase) tend to react faster with nitric oxide than heme proteins that have all six coordination positions occupied (i.e. cytochrome c), as nitric oxide binding in the latter requires displacement of the 6th ligand, usually an amino acid. In the case of heme proteins which have an accessible 6th coordination position such as hemoglobin, myoglobin, and guanylate cyclase, the presence of a histidine coordinated to the heme stabilizes the binding of water at the 6th coordination position of the ferric state, making it more difficult for nitric oxide to bind to the heme, in comparison with the ferrous form (11). The presence of the protein backbone makes the combination reaction of heme proteins with nitric oxide slower than with free heme, but more selective, since the relative apolar environment present in the heme cleft makes diffusion of NO much more favorable than that of NO-derived oxides such as nitrite. In addition, dissociation of the nitrosyl-heme protein complex is usually slower than in the model heme. In Table 1 the set of constant values (kf, kd, and Keq) for ferrous and ferric myoglobin and cytochrome c are shown as representative examples of the general characteristics of nitric oxide binding to heme proteins (12, 13). Association-dissociation of NO in heme proteins have implications for the biological role of NO. For instance, the termination of the stimulation of guanylate cyclase by NO relies on the dissociation of NO from the heme2+ of the enzyme. c. Formation of Heme-Nitrosyl Complexes. The formation of complexes between NO and heme can be detected both by optical and electron paramagnetic resonance studies (5-10, 14). The formation of an NOheme2+ complex results in an EPR signal at g ) 2.0, indicating a iron complex with rhombic symmetry (5, 7-9). On the other hand, the NO-heme3+ complex is diamagnetic and therefore EPR silent. The interactions of NO with heme may lead to the formation of either a stable or transient heme-nitrosyl complexes. In some cases, the heme-nitrosyl is a stable complex in equilibrium with free NO and does not result in net redox change at the iron. In other cases, the formation of a transient complex is an intermediate step during NO-heme interactions that lead to redox changes at the metal center and formation of nitrite or nitrate. Finally, there are examples in which the initial formation of a relatively stable heme-nitrosyl complex is followed by a slow redox process. Examples of the different stabilities of heme-nitrosyl complexes are found during the interactions between NO and hemoglobin and myoglobin and will be discussed in section IId. The formation of heme-nitrosyl complexes can modify the reactivity of NO and make it more available for
Table 1. Rate and Equilibrium Constants for Nitric Oxide Reactions with Heme Proteinsa heme protein
kf (M-1 s-1)
kd (s-1)
Keq (M-1)
Mb2+ Mb3+ Cyt c2+ Cyt c3+
1.7 × 107 1.9 × 105 8.3 7.2 × 102
1.2 × 10-4 1.4 × 101 2.9 × 10-5 4.4 × 10-2
1.4 × 1011 1.4 × 104 2.9 × 105 1.6 × 104
a Data for myoglobin (from whale) and cytochrome c (from horse) were obtained from refs 12 and 13. Reactions were carried out at pH 6.5-7.0 and T ) 20 °C.
Table 2. Reduction Potential of Some Iron Complexes redox couple
E°′ (V)
Fe3+/Fe2+ Fe3+-EDTA/Fe2+-EDTA Fe3+-DTPA/Fe2+-DTPA Fe3+-hemoglobin/Fe2+-hemoglobin [Fe-O2]2+-hemoglobin/Fe3+-hemoglobin + H2O2 Fe3+-myoglobin/Fe2+-myoglobin [Fe-O2]2+-myoglobin/Fe3+-myoglobin + H2O2 Fe3+-cytochrome c/Fe2+-cytochrome c Fe3+-cytochrome a3/Fe2+-cytochrome a3
0.77 0.12 0.03 0.14 0.30 0.05 0.39 0.27 0.55
nitrosylation reactions (15). The reaction of NO with methemoglobin (Hb3+) (or metmyoglobin) (eq 3) results in the formation of a complex which either can be hydrolyzed to deoxyhemoglobin (Hb2+) and nitrite (eq 4) or can transfer the nitrosyl group to nucleophilic compounds (i.e., thiols, eq 5).
Hb3+ + NO h [Hb3+...NO] f Hb2+-NO+
(3)
Hb2+-NO+ + H2O f Hb2+ + NO2- + 2H+
(4)
Hb2+-NO+ + RS- f Hb2+ + RSNO
(5)
In the first step, nitric oxide binds to the ferric heme and transfers an electron to the iron. The resulting heme2+NO+ complex constitutes a nitrosylating compound due to the NO+ character of NO in the complex, and the overall two-step sequence can be described as a reductive nitrosylation process. However, there is diversity in the stability of NOheme3+ complexes. While the complexes of ferric heme and myoglobin are unstable and gradually transform to ferrous-NO+ complexes which tend to react and nitrosylate, the complexes with the ferric forms of horseradish peroxidase, cytochrome c peroxidase, and cytochrome c are quite stable (1, 13, 16-18). In this regard, electron transfer from NO to the heme depends on the redox potential of the iron and is not much influenced by the electron withdrawing character of the porphyrin periphery (1). In turn, the redox potential of the iron complex in heme proteins, which in all cases is significantly lower than that of free iron (Table 2), will be detemined by the combination of the nature of the porphyrin with its substituents as well as the amino acids linked to the 5th and 6th coordination position of the heme, which influences the electron density of the iron (19). d. Reactions of NO with Hemoglobin. We will describe these reactions in more detail as they are examples of different types of interactions that NO may have with heme proteins. NO reacts with oxyhemoglobin (Hb2+-O2) at 3 × 107 M-1 s-1 (20) leading to the formation of met-Hb and nitrate (eq 6). It has been postulated that this reaction occurs with the intermediate formation of peroxynitrite anion. NO can also react with methemoglobin leading to deoxyhemoglobin and a nitrosylating species (eq 7). Finally, the reaction of NO with deoxyhemoglobin with
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Scheme 1
rate constants above 107 M-1 s-1 (Table 1) leads to the formation of a stable nitrosylhemoglobin complex (Hb2+NO).
Hb2+-O2 + NO S Hb3+ + NO3-
(6)
Hb3+ + NO f Hb2+-NO+
(7)
Hb2+-NO+ + H2O f Hb2+ + NO2- + 2H+
(8)
Hb2+ + NO f Hb2+-NO
(9)
Thus, oxyhemoglobin reactions with excess NO will ultimately yield the stable hemoglobin-nitrosyl complex, through a sequence of reactions as shown in Scheme 1. The reactions of NO with oxyhemoglobin in the intravascular space will mostly lead to the formation of methemoglobin (eq 6) due to the large excess of oxyhemoglobin over NO in red blood cells. This reaction represents a drain for NO produced in tissues and keeps NO with the characteristics of a local signal transduction molecule. At low oxygen tensions, where part of the ferrous hemoglobin will be in the deoxygenated form, formation of nitrosylhemoglobin can occur. e. NO Inhibition of Cytochromes-Catalyzed Reactions. Of particular interest in relation to cytotoxic effects of nitric oxide are the interactions of nitric oxide with cytochrome P450 and cytochrome c oxidase (cytochrome aa3). Nitric oxide interactions with cytochrome P450 result in inhibition of enzymic activities toward organic substrates. Reversible nitric oxide binding to cytochrome P450 heme prevents oxygen binding and monooxygenase activity (21, 22). For cytochrome c oxidase, a similar inhibitory mechanism involving NO binding to the sixth coordination position of the heme of cytochrome a3 has been shown (18-23). Since oxygen is the terminal electron acceptor of the respiratory chain, NO binding to the heme results in inhibition of mitochondrial electron transport (24- 26). The efficiency of NO binding to cytochrome a3 depends on oxygen tension since both gaseous ligands compete for the same binding site. Thus, NO will cause a more profound and persistent inhibition of cellular respiration at lower oxygen tensions (26). In any case, the binding constant for NO is significantly larger than that for oxygen (25).
III. Iron-Sulfur Cluster Proteins: The Example of Aconitase NO has been shown to bind to iron-sulfur clusters leading to the formation of iron-nitrosyl complexes (for
Figure 1. Inhibition of aconitase activity by nitric oxide. Aconitase (30 µM) was anaerobically incubated without (+) or with 50 (2), 100 (9), and 200 (b) µM nitric oxide in 100 mM Tris-HCl, pH 7.8, 25 °C. The arrow shows the time when an argon stream was introduced into the 200 µM nitric oxide reaction mixture. Argon was flushed for 5 min, and recovered activity is represented with open circles. (From ref 36.)
a review, see ref 27). NO, having an unpaired electron, binds to Fe2+ forming an EPR active complex with a characteristic EPR spectrum. These iron-nitrosyl complexes may be stable and reversible, or nitric oxide may progressively lead to cluster oxidation and disruption. A key interaction of NO with metalloproteins in regard to cytotoxicity is represented by the reaction of NO with mitochondrial aconitase (28, 29). Mitochondrial aconitase is an iron-sulfur containing enzyme that participates in the Krebs cycle. The iron-sulfur cluster of aconitase is of the type [4Fe-4S] and has the peculiarity of having one of the four atoms of iron with a free uncoordinated position to sulfur. For aconitase and other dehydratases that share similarities in the iron-sulfur cluster structure, the fourth iron atom, also known as the iron-R (FeR), is a labile iron, that can be attacked by different oxidizing electrophiles such as molecular oxygen, ferricyanide, and superoxide radical, with the concomitant release of the FeR and cluster disruption (30-34):
[4Fe-4S]2+ f [4Fe-4S]3+ + e-
(10)
[4Fe-4S]3+ f [3Fe-4S]1+ + Fe2+
(11)
The disruption of the [4Fe-4S] cluster leads to enzyme inactivation, and it has been largely assumed in the literature that the interaction of macrophage-derived nitric oxide with target cells, leading to the inactivation of aconitase, was due to direct reactions of nitric oxide with the cluster leading to the formation of the [3Fe-4S] inactive form. Recent experiments by Hausladen and Fridovich (35) and ourselves (36), however, indicate that nitric oxide tends to form a relatively stable reversible iron-nitrosyl complex with aconitase. Iron-nitrosyl complex formation leads to transient inhibition of aconitase activity but does not cause cluster disruption and inactivation as readily as previously thought (Figure 1). Still, it is possible that, after formation of the ironnitrosyl complex, nitric oxide could slowly oxidize the cluster. On the other hand, peroxynitrite2 causes cluster 2The term peroxynitrite refers to both peroxynitrite anion (ONOO-) and peroxynitrous acid (ONOOH) throughout the text.
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Table 3. Inactivation of Aconitase by Peroxnitrite and Reactivation by Thiols and Ferrous Irona condition
% initial activity
control ONOOROA ONOO+Cys +DTT +GSH +Fe2+ +Fe2+, Cys +Fe2+, DTT +Fe2+, GSH
100 ( 9 31 ( 9 103 ( 15 71 ( 2 40 ( 3 46 ( 4 41 ( 6 99 ( 0 109 ( 7 60 ( 1
a Aconitase (30 µM) was exposed to 100 µM peroxynitrite in 100 mM Tris-HCl, pH 7.8, for 5 min at 25 °C. ROA is a “reverse order addition” experiment in which peroxynitrite first decomposes in buffer alone and the enzyme is added 5 min after. Then, thiols (10 mM) and ferrous iron (100 µM) were added and incubated with the enzyme for 5 min and enzyme activites assayed immediately after (from ref 36).
disruption and inactivation of aconitase (Table 3).
[4Fe-4S]2+ + ONOOH f [3Fe-4S]1+ + Fe3+ +OH- + NO2- (12) Peroxynitrite-mediated cluster disruption did not go beyond the [3Fe-4S] state, since post-incubation with ferrous iron and thiols resulted in cluster reassembly and recovery of enzymic activity (Table 3). Interestingly, it has been recently demonstrated that cytosolic aconitase is an iron-responsive element-binding protein (IRE-BP) that requires complete disassembly of the cluster to carry out its function (37, 38). In this regard, NO and NO-derived products have been proposed to participate in regulation of the activity of the IRE-BP activity of cytosolic aconitase. After initial cluster disruption to [3Fe-4S] by nitrogen oxides, a follow-up sequence to complete cluster disruption is required. Biologically relevant concentrations of either NO or even peroxynitrite are unable to continue cluster disruption beyond the [3Fe-4S], and for this reason it is possible that an enzymic mechanism may participate to complete the disruption process. Although NO itself does not appear to be very efficient in aconitase inactivation, nitrosothiols, typically used as NO donors in biochemical systems, can efficiently cause aconitase inactivation (Figure 2). The extent of aconitase inactivation is not related to the exposure of the enzyme to the levels of free NO released by the nitrosothiols but instead is related to the nitrosothiol properties and concentration (36). The most likely mechanism for nitrosothiol-mediated inactivation of aconitase is the reaction of aconitase with the -NO moiety of the nitrosothiols which has a character of nitrosonium (NO+), a strong electrophilic agent and much stronger oxidant than free NO (4). The reaction of the NO+ moiety of the nitrosothiols would be responsible for cluster disruption. These observations also indicate that interpretation of results obtained using chemical NO donors must take into consideration the possibility of direct reactions between the NO donor and the target molecule under study.
IV. NO-Ferritin Interactions NO-mediated disruption of iron homeostasis has been postulated for many years as one important cytotoxic mechanism. Indeed, target cells lose a major fraction of cell and mitochondrial iron during coincubation with NOproducing biochemical or cellular systems.
Figure 2. Aconitase inactivation by nitrosoglutathione. Aconitase (30 µM) was anaerobically incubated with 0 (O), 5 (2, b) or 10 (() mM GSNO in 100 mM Tris-HCl, pH 7.8, at 25 °C. The arrow indicates a second addition of 5 mM GSNO. (From ref 36.)
Ferritin is an iron storage protein containing up to 4500 atoms of iron per molecule, mostly in the ferric form. Most of the cellular non-heme iron is stored as ferritin. Since the EPR spectra of liver cells and of ferritin in the presence of NO are strikingly similar, it appears reasonable to postulate that ferritin participates in the formation of non-heme iron-nitrosyl complexes in cells. In mammalian ferritins, three types of EPR-distinguishable iron-nitrosyl complexes have been described (39). The EPR signals (A, B, and C) have been attributed to complexes between ferrous iron and nitric oxide at imidazole groups of histidine, thiols groups of cysteines, and carboxylate groups of aspartate and glutamate, respectively. The biological significance of the formation of iron-nitrosyl complexes in ferritin is unknown at present, but it has been speculated that nitric oxide may be stored in cells by this mechanism (39). Apart from the formation of reversible complexes between Fe2+ and nitric oxide, nitric oxide may also mediate release of stored Fe3+ from ferritin (40). The molecular mechanisms by which NO can promote iron mobilization from ferritin have not been described, but nitric oxide may reduce ferric to ferrous iron and that complex may slowly dissociate with the concomitant release of iron in the reduced form. In this regard, it has been previously shown that superoxide radical anion (O2-) mediates iron reduction and release from ferritin (41).
V. Evidence for the in Vivo Interactions between NO and Iron-Containing Proteins There are two lines of evidence that demonstrate the interaction between NO and iron-containing proteins in living systems. The first line of evidence relies on the formation of iron-nitrosyl complexes in cells and tissues (42-45). For example, early experiments (42) showed that vegetative cells of Clostridium botulinum incubated in the presence of an NO-generating system (ascorbate plus nitrite) led to the formation of cellular iron-nitrosyl complexes evidenced by the appearance of the characteristic EPR signal at g ) 2.035 (77 K). Later experiments showed that exposure of cells to authentic nitric oxide or even nitric oxide released by activated cytotoxic macrophages resulted in the formation of iron-nitrosyl
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complexes in the target cell (43-45). The formation of these complexes has been postulated to participate in the cytotoxicity associated with NO. The second line of evidence is the formation of the stable nitrosylhemoglobin complex in tissues that have been exposed to enhanced endogenous or exogenous fluxes of NO (46-48). For example, nitrosylhemoglobin formation, as detected by EPR, is increased in different pathological situations where there is an overproduction of NO (46, 47). In this case, nitrosylhemoglobin is used as a “footprint” of nitric oxide production in vivo. The formation of nitrosylhemoglobin is favored in tissue compartments which have low oxygen tension as oxyhemoglobin will evolve to methemoglobin during reactions with NO. Also, NO exposures to animals lead to an increased concentration of nitrosyl- and methemoglobin in blood (48).
VI. NO Interactions with Zinc-Sulfur Clusters Zinc is the second most abundant transition metal in higher animals, and zinc in proteins is usually complexed with protein ligands in a structure known as “zinc finger” domains. These domains consist of zinc bound to either the thiol group of cysteine and/or the imidazole nitrogens of histidine. For some proteins such as alcohol dehydrogenase, the zinc finger is critical for catalytic activity; in the case of zinc-sulfur-containing transcription factors, zinc fingers are essential for specific DNA binding. In one report (49), it has been shown that the zinccomplexing protein metallothionein, which has 20 cysteines per molecule and complexes up to 7 Zn2+, interacts with NO leading to transient S-nitrosylation of the thiol moieties followed by disulfide formation and Zn release. There is a good correlation between thiol oxidation and Zn release. The same authors showed that NO treatment of a zinc-sulfur cluster-containing transcription factor leads to inhibition of DNA-binding activity. NO-mediated disruption of Zn-sulfur clusters may represent a mechanism of cytotoxicity and cytostasis that requires further investigation. Still, the precise mechanisms by which NO can disrupt the cluster are not clear. NO itself does not oxidize thiol groups at the physiological pH range. NO-mediated oxidation and disruption of zinc-sulfur clusters may require the intermediate formation of nitrogen dioxide by the reaction of nitric oxide with molecular oxygen or peroxynitrite by the reaction of NO with superoxide. Both nitrogen dioxide and peroxynitrite are strong sulfhydryl oxidizing agents. NO exposures under strict anaerobic conditions can rule out the contributory role of nitrogen dioxide as well as the potential formation of peroxynitrite which may arise from superoxide formation secondary to thiol autoxidation.
VII. NO Interactions with Copper-Containing Proteins Different copper-containing proteins including hemocyanin, tyrosinase, and blue copper oxidases (ceruloplasmin, ascorbate oxidase, and laccase) have been shown to react with NO (50-52). In general terms, the reactions can result in either formation of stable copper-nitrosyl complexes or redox chemistry. However, in many cases it has been difficult to discriminate between both mechanisms as reduction of the copper or complex formation may lead to the same or similar EPR and optical properties.
Radi
The reactivity of NO with copper-containing proteins is largely dictated by the structural and geometric characteristics of the copper-protein complex at the active site. In this regard, three types of copper (1, 2, and 3) have been described in copper-containing proteins, with some proteins having more than one type of copper in its structure. NO reacts with copper proteins which have types 1 and 3 copper but not with type 2. Copper proteins containing type 1 copper (i.e., plastocyanin, azurin, stellacyanin) have the copper ligated to two nitrogen atoms of histidines, one sulfur from a methionine, and one sulfur from a cysteine. NO predominantly reacts with type 1 copper proteins in the oxidized state with the formation of stable coppernitrosyl complexes. The initial combination reaction of NO with oxidized copper (Cu2+) is followed by copper reduction and formation of a Cu1+-NO+ complex (eq 13). This complex is photodissociable at 77 K. Coppernitrosyl complex formation leads to disappearance of the typical EPR signal of type 1 copper and of the 600 nm optical absorption band.
Cu2+ + NO h Cu1+-NO+
(13)
Type 3 copper-containing proteins (i.e., hemocyanin, tyrosinase, blue oxidases) contain one or more dinuclear copper sites which are EPR silent in both the reduced Cu1+ Cu1+ and oxidized Cu2+ Cu2+ form. Of the two coppers of the dinuclear site, Cu(B) is liganded by three histidines, which are located in a 56 amino acid highly conserved sequence homology among the different type 3 copper-containing proteins. In contrast, the structure accounting for Cu(A) binding is less conserved, although in general two or three histidines are involved as Cu(A) ligands. Type 3 copper proteins react with nitric oxide in both the oxidized and reduced state to form nonphotodissociable copper-nitrosyl complexes, which can be followed by redox chemistry with disruption of the complex. Type 2 copper is the most abundant in nature, but, in contrast to type 1 and 3 copper, does not form a homogeneous group, being very variable in structure, function, and properties. Although not possible to define a common coordination structure, their geometry is usually square planar. Type 2 copper proteins do not react with nitric oxide. For instance, the cytosolic form of the enzyme superoxide dismutase contains a type 2 copper in the active site which is liganded to histidines and does not react with nitric oxide. Indeed, neither the optical or the EPR spectrum of SOD is affected by nitric oxide (50) nor is nitric oxide significantly consumed in the presence of SOD.3 An interesting protein to analyze in the context of a potential cytotoxic action of NO is cytochrome c oxidase, known to have two coppers, CuA and CuB, in its structure. While CuA does not react with NO, resembling type 2 copper, CuB binds to NO, leading to the formation of photodissociable copper-nitrosyl complexes as those described for type 1 copper (53). Therefore, NO can bind at two different metal centers in cytochrome c oxidase, the heme and the CuB of cytochrome a3. While the biological consequence of NO binding to the heme is better understood, the effect that the formation of the NO-CuB complex may have in electron transport is unknown at present. 3R.
Radi, unpublished observations.
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Scheme 2
Table 4. Rate Constants for Peroxynitrite Reactions with Metalloproteins
Nitric Oxide and Superoxide Dismutase. Interestingly, superoxide dismutase, one of the most abundant copper proteins in nature and present in eukaryotic cytosols, does not react with NO, but it has an enormous influence in the biological half-life and fate of NO. How is this? Superoxide dismutase increases the half-life of NO because it minimizes the almost diffusion-controlled reactions of nitric oxide with superoxide (54-57). Superoxide is continuously produced in aerobic cells, and its reaction with NO not only decreases its availability but also leads to the formation of the highly oxidizing and cytotoxic species, peroxynitrite anion (58-61). Alternatively, it has been proposed that superoxide dismutase can enhance the concentration of nitric oxide in living systems via a one-electron oxidation of nitroxyl anion (62).
Peroxynitrite is formed by the fast reaction (k ) 6.7 × 109 M-1 s-1; 57) between superoxide and NO and has a half-life of less than a second at pH 7.4 due to its rapid proton-catalyzed (pKa ) 6.8; 58) decomposition to an oxidant with reactivity similar to that of hydroxyl radical and nitrogen dioxide.
(14)
ONOO + H h ONOOH f “ OH” + NO2 (15) -
+
•
k (M-1 s-1)
ref
myeloperoxidase HRP alcohol dehydrogenase aconitase cytochrome c2+ SOD ceruloplasmin Fe3+-EDTA
2 × 107 3 × 106 4 × 105 2 × 105 2 × 105 103-105 NDa 6 × 103
65 65 66 36 67 54 68 69
ND ) not determined.
lead to disruption of the metal centers with subsequent metal release. In the case of aconitase and alcohol dehydrogenase, the disruption of the clusters leads to loss in enzymic activities (36, 66). Peroxynitrite also reacts with the three different classes of superoxide dismutases (Cu/Zn-, Mn-, and Fecontaining), leading to the formation of nitrating intermediates. While peroxynitrite does not significantly inactivate Cu/Zn SOD, it does inactivate the Mn and Fe forms (64). Peroxynitrite-mediated SOD-catalyzed nitration of critical tyrosine residues in proteins has been postulated to participate in cellular degeneration and apoptosis (70-72).
IX. Nitric Oxide Modulation of Free Radical Processes
VIII. Peroxynitrite Interactions with Metalloproteins
O2•- + •NO f ONOO-
a
reaction
•
Peroxynitrite is a much stronger oxidant than any of its precursors. Indeed, the one-electron redox potential of the ground state form of peroxynitrite has been estimated to be +1.4 V and of the energized intermediate (ONOOH*) about +2.1 V, close to the couple hydroxyl radical/water which is +2.3 V (4). Peroxynitrite can also promote two-electron oxidations with the ground state form (63), which has a two-electron redox potential of +0.9 V (4) (Scheme 2). Due to its instability and reactivity, the reactions of peroxynitrite with metalloproteins are oxidation reactions, without formation of stable reversible complexes. In some cases the reactions of peroxynitrite with metal centers in proteins lead to protein nitration via formation of a nitronium-like (NO2+) species (64, 65). Different metalloproteins have been shown to interact with peroxynitrite at significant rates (Table 4). The second order rate constants for the reported reactions of the ground state form of peroxynitrite with metalloproteins range from 103-107 M-1 s-1 (36, 65-69) and represent the fastest studied reactions for peroxynitrite up to now. The reactions in some cases lead to a reversible redox change of the metal center (i.e., peroxidases and cytochrome c; 65, 67). In other cases, such as in the case of the reactions of peroxynitrite with ironsulfur cluster containing aconitase (36), zinc-sulfur cluster containing alcohol dehydrogenase (66), and the copper storage protein ceruloplasmin (68), the reactions
It is well established that superoxide radical anion and hydrogen peroxide, byproducts of the partial reduction of molecular oxygen in biological systems, participate in oxidation and biomolecular damage (73-75). Although superoxide and hydrogen peroxide can directly attack a limited number of targets, an important part of their toxicity relies on the formation of the highly oxidizing hydroxyl radical, the strongest oxidant produced in biology (E°′ ) +2.3 V; 76). Formation of hydroxyl radical requires the participation of a suitable transition metal such as iron or copper, in a process known as the HaberWeiss reaction.
O2•- + Fe3+ f O2 + Fe2+
(16)
Fe2+ + H2O2 f •OH + OH- + Fe3+
(17)
In many cases the transition metal is provided as a low molecular weight complex, but also some metalloproteins can provide the metal. For instance, hemoglobin has been postulated as a Fenton reagent (77). Metalloproteins, including heme proteins in the ferric form (hemoglobin, myoglobin, and cytochromes), can react with hydrogen peroxide leading to the formation of higher oxidation states of iron, forming iron-oxo complexes such as ferryl (Fe4+dO) (78). Similar intermediates are found during the catalytic action of heme peroxidases and catalase with the formation of the “compound I” of peroxidases. The iron-oxo complexes have a reactivity similar to that of hydroxyl radical, and it has been shown that metalloprotein interactions with reactive oxygen intermediates lead to oxidation reactions. Also, in many heme proteins, formation of higher oxidation states of iron leads to a progressive destruction of the heme followed by iron release, which may promote HaberWeiss chemistry (79). Nitric oxide inhibits these processes by binding to ferrous or ferric iron, precluding its interaction with hydrogen peroxide (80, 81). As an additional mechanism to “downregulate” free radical chemistry, nitric oxide can reduce ferryl iron back to the
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Chem. Res. Toxicol., Vol. 9, No. 5, 1996 Scheme 3
ferric state, inhibiting the reaction of the high oxidation state iron with other more critical moieties (82). In the case of cytochrome P450 (21) and some oxygenases such as lipoxygenase, nitric oxide binds to the ferrous form and inhibits further redox/free radical chemistry (83). Lipoxygenase is a non-heme non-ironsulfur-containing iron enzyme which participates in arachidonic acid metabolism and the formation of lipid hydroperoxides. Nitric oxide inhibits lipooxygenasemediated lipid oxidation by binding to the ferrous form of the enzyme, with fatty acid binding to the enzyme being competitive to that of nitric oxide (83). In Scheme 3, the antioxidant mechanisms of nitric oxide during its interaction with heme proteins are shown in steps a, c, and e which inhibit pro-oxidant processes promoted in steps b and d. The inhibitory role that nitric oxide appears to play in metalloprotein interactions with reactive oxygen species, therefore, resulting in a “antioxidant” effect, may be counterbalanced by two other actions that nitric oxide may promote such as iron mobilization from ferritin and activation of the cytosolic aconitase (IRE-BP), both processes resulting in an increased concentration of “free” iron available to carry out Haber-Weiss and Fenton reactions. In addition, it has recently been shown that induction of nitric oxide synthesis in isolated rat hepatocytes results in near 60% loss of catalase activity, a key hydrogen peroxide detoxifying enzyme (84). The elucidation of the complex interrelationships between pro- and antioxidant actions of nitric oxide in the context of its interactions with metalloproteins and cytotoxicity will require important research efforts in the next few years.
Acknowledgment. I thank Drs. Ana Denicola, Gerardo Ferrer, Harry Ischiropoulos, Willem H. Koppenol, and Homero Rubbo for critical reading of the manuscript. This work was supported by grants from SAREC (Sweden) and CONICYT (Uruguay).
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