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Chem. Res. Toxicol. 1996, 9, 836-844
Oxidative Damage and Tyrosine Nitration from Peroxynitrite Joseph S. Beckman* Departments of Anesthesiology and Biochemistry, The University of Alabama at Birmingham, Birmingham, Alabama 35233 Received August 21, 1995
The most widely accepted and taught mechanism of free radical toxicity at present is the formation of hydroxyl radical by the superoxide-driven Fenton reaction. The reaction requires two steps, where superoxide first reduces ferric iron to ferrous iron: 106 M-1‚s-1
Fe3+ + O2•- 98 Fe2+ + O2 and then the ferrous iron attacks hydrogen peroxide to generate hydroxyl radical (1): 103-105 M-1‚s-1
Fe2+ + H2O2 98 Fe3+ + HO• + HOHydroxyl radical is highly reactive and certainly capable of destroying isolated DNA, protein, or lipid in a test tube. However, a highly reactive species is not necessarily highly toxic. In the case of hydroxyl radical, the rate of reaction with every organic molecule is near the diffusion limit, ranging from 109 to 1010 M-1‚s-1. Thus, hydroxyl radical will randomly damage isolated DNA in a simple phosphate-buffered solution. Binding of DNA binding proteins protects sequences from hydroxyl radical, enabling molecular biologists to identify specific regions of the DNA that interact with the protein. As more and more components of a cell are added to the mixture, the chances that DNA or any critical cellular target will be hit by a highly reactive species become smaller and smaller. Radiation chemists calculate that the diffusion distance of hydroxyl radical is only 3 nm in a cell (2). Hydroxyl radical will randomly attack noncritical cellular components because its broad reactivity at diffusionlimited rates makes hydroxyl radical an indiscriminant species. The reductionist nature of biochemistry to work on clean, isolated systems exaggerates the apparent toxicity of hydroxyl radical because alternative noncritical targets are removed from the system. An additional difficulty is that natural antioxidants and scavenging enzymes like superoxide dismutase and catalase are absent from in vitro systems but are present in vivo. Still, the injection of exogenous superoxide dismutase or catalase into animals and humans can reduce injury from ischemia and inflammation, which establishes that oxidative processes are involved in many disease processes that cannot be prevented by endogenous antioxidant defenses (3, 4). However, the superoxide-driven Fenton reaction is a relatively slow reaction that is easily stopped in vitro by adding small amounts of catalase or superoxide dismutase. In vivo, up to 1% of soluble protein is superoxide dismutase (5, 6), while catalase and glutathione peroxidase are also abundant and effectively remove hydrogen peroxide. Even submicromolar concentrations of hydrogen peroxide will not * Tel: 205-934-5422; FAX: MS.THT.ANES.UAB.EDU.
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exist for long in vivo. Yet, millimolar concentrations are called low doses in most in vitro experiments. To understand oxidant toxicity in vivo, one needs to look for free radical reactions that are fast enough to outcompete endogenous antioxidant defenses. The reactions must be fast enough that neither superoxide dismutase nor catalase will stop the reaction. Other oxidative mechanisms must be operating in vivo. In the present review, we will focus upon only one reactionsthe diffusion-limited reaction of superoxide with nitric oxide. More certainly remain to be investigated. At present, the only known biological molecule that is produced in high enough concentrations and can react fast enough with superoxide to outcompete endogenous superoxide dismutase is nitric oxide (6). Nitric oxide is a free radical that combines by radical-radical coupling with superoxide to form peroxynitrite anion (7):
O2•- + •NO f -OONO The reaction rate is 6.7 × 109 M-1‚s-1, which is at least 3 times faster than superoxide dismutase reacts with superoxide (8, 9). The scavenging of superoxide by superoxide dismutase is partially reduced by physiological levels of chloride ions, which screen the electrostatic field that attracts superoxide to the active site. Chloride decreases the scavenging of superoxide by a factor of 2-3 (9, 10), but should have no effect upon the radical-radical coupling of nitric oxide and superoxide. Therefore, the formation of peroxynitrite is slightly more favorable under physiological conditions compared to reactions conducted in phosphate buffer. Peroxynitrite itself is not a free radical because the two unpaired electrons on superoxide and nitric oxide have combined to form a new bond. Peroxynitrite is an isomer of nitrate, but is 36 kcal‚mol-1 higher in energy (11). Peroxynitrite is relatively stable in alkaline solution and can be stored in millimolar concentrations at -80 °C for months (12). It has been prepared in a pure solid form as the tetramethylammonium salt, which is indefinitely stable (13). Irradiation of crystals of potassium or rubidium nitrate with short wavelength UV light causes the crystals to turn yellow (14, 15). The yellow color results from the accumulation of 1-3% peroxynitrite in the crystal matrix, which remains stable, when protected from light, for at least 3 years. The formation of peroxynitrite in the Martian soil by UV irradiation of nitrate may account for the evolution of oxygen and decarboxylation of amino acids discovered during the Viking Mars missions (16). The unusual stability of peroxynitrite results from the anion being folded in the cis conformation (Figure 1). Our initial evidence for the cis isomer being predominant in solution was based upon the Raman spectra (17). More recently, we have found that the 15N NMR shifts and the UV spectra are also more consistent with peroxynitrite © 1996 American Chemical Society
Forum on Nitric Oxide: Chemical Events in Toxicity
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and hydrogen peroxide, which was first investigated by Bayer in 1902 (15). The formation of hydroxyl radical from peroxynitrous acid in acidic solutions of nitrite and hydrogen peroxide was proposed by Halfpenny and Robinson (23, 24) as well as Mahoney (25). We investigated the formation of a hydroxyl radical-like oxidant from peroxynitrite at neutral pH for the oxidation of dimethyl sulfoxide and deoxyribose.
ONOO- + H+ S ONOOH f “HO• + •NO2” f NO3-
Figure 1. Structures of cis, trans peroxynitrite, and nitrate. The two oxygens in cis peroxynitrite interact weakly, which helps stabilize peroxynitrite. However, cis peroxynitrite cannot directly rearrange to peroxynitrite. trans peroxynitrite can directly isomerize to nitrate by the terminal oxygen swinging out to attack the nitrogen. Consequently the decay of peroxynitrite depends upon cis peroxynitrite becoming protonated, isomerizing to trans peroxynitrous acid, and then rearranging to nitrate. The activated “HO•” state might be a triplet rather than singlet state.
being in the cis form (18-20). Ab initio calculations that take electron correlation into account predict the cis anion is about 3.5 kcal‚mol-1 more stable than the trans anion (17). In the cis conformation, the negative charge is partially delocalized over all four atoms, with a significant interaction between the two terminal peroxide oxygens. The terminal oxygen of peroxynitrite cannot directly approach the nitrogen to form nitrate without substantial rearrangement of the other oxygens, which contributes to the unusual stability of peroxynitrite. The terminal peroxide oxygen on the trans anion can directly approach the nitrogen by a stretching of the O-O bond and bending of the N-O-O bond. The N-O bond in the middle of cis peroxynitrite is represented as a single bond, but the bond order is closer to 1.5 (15). Consequently, cis and trans peroxynitrite cannot directly interconvert, because there is a significant energy barrier for isomerization between the cis and trans anions, estimated to be 30 kcal‚mol-1 by ab initio calculations. Because of the high barrier for interconversion between the cis and trans anion, the conformation of peroxynitrite and the pH of the solution can dramatically alter its reactivity and affect its toxicity (12, 21). Peroxynitrite exhibits a pKa at 6.8 in phosphate buffer, and when protonated, the barrier for cis-trans interconversion is reduced to 13 kcal‚mol-1 based upon gas phase calculations (22). In effect, the hydrogen ion is neutralizing the negative charge that is partially delocalized over the anion. However, the barrier for cis-trans isomerization of the acid may be a few kilocalories higher because gas phase calculations do not take into account hydrogen bonding to water. We have proposed that the barrier for acid isomerization is slightly higher and may be rate limiting under some circumstances in the reaction of peroxynitrite with some molecules. Perhaps the most important example is in the reaction of peroxynitrite with Cu,Zn superoxide dismutase that is described in greater detail below. Peroxynitrous acid has long been known to be a strong oxidant. It is formed by the reaction of acidified nitrite
The hydroxyl radical and nitrogen dioxide are written in quotes because the O-O bond of peroxynitrous acid apparently does not break homolytically to give free hydroxyl radical and free nitrogen dioxide. Rather, peroxynitrous acid seems to form a highly reactive intermediate with the reactivity of hydroxyl radical. There are several lines of evidence for concluding that hydroxyl radical is not formed as a free species. Hydroxyl radical and nitrogen dioxide react to form peroxynitrite and not nitrate (26). The activation entropy for peroxynitrite decomposition is small, which suggests that the transition state is rather rigid and symmetric, whereas homolytic fission of a bond has much higher entropies. To account for these observations, we proposed that peroxynitrite in the trans configuration could become vibrationally activated to form a species with the reactivity of hydroxyl radical (Figure 1) (11). The two vibrations were stretching of the O-O bond and bending of the N-O-O bond angle, the same motions that are necessary for peroxynitrite to rearrange to nitrate. Recent calculations by T. Hamilton and G. Tsai (Department of Chemistry, University of Alabama at Birmingham) indicated that there is a low lying triplet state (27), which has confounded low level calculations on peroxynitrite. The availability of a low lying triplet state suggests that peroxynitrite may effectively form a biradical (Figure 1), which contributes to its ability to initiate free radical reactions even though peroxynitrite is not a radical itself. Curiously, the oxidation yield from the hydroxyl radical-like pathway decreased at slightly alkaline pH, exhibiting apparent pKa’s at 7.6 and 7.8 for deoxyribose and DMSO, respectively (28). These pKa’s were significantly higher than the pKa of peroxynitrite estimated in phosphate buffer (11, 29, 30). Subsequently, we found that the formation of nitrogen dioxide from the reaction of peroxynitrite also decreased at alkaline pH with an apparent pKa of 7.8 (31). To explain the loss of oxidation potential at alkaline pH, we proposed that both the trans acid and anion can decompose to form nitrate (21, 31). However, the trans acid is far more likely to act as an oxidant and attack other molecules, because the anion already has one negative charge. The anion is unlikely to extract a second electron to become doubly charged. When the pH is higher than 8, trans peroxynitrous acid will ionize to the anion and decompose to nitrate without acting as a strong oxidant, accounting for the loss of reactivity at alkaline pH (Figure 2). The difference in the apparent pKa’s between the decomposition of peroxynitrite in phosphate buffer and the decrease in reactivity may be due to a slight difference in the pKa’s of the cis and the trans forms of peroxynitrite (17). The cis form of peroxynitrite anion is estimated to be approximately 3 kcal‚mol-1 lower in energy than the trans anion. The energy difference between the two acid conformations is only 1 kcal‚mol-1. Consequently, the energy for trans peroxynitrous acid to
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Figure 2. Decomposition of trans peroxynitrous acid and anion to nitrate through an activated intermediate. Once cis peroxynitrous acid has isomerized to trans peroxynitrous acid, trans peroxynitrous acid will be in rapid equilibrium with its conjugate base as determined by the pH. As trans peroxynitrous acid rearranges to nitrate, it forms a high energy intermediate that is proposed to have the reactivity of hydroxyl radical (11). The trans anion will also rearrange to nitrate, but is unlikely to be an oxidant because it is already negatively charged. It is also possible that the activated intermediate with hydroxyl radicallike reactivity is a triplet state, where a lone electron pair is split to occupy two orbitals.
ionize should be 2 kcal‚mol-1 higher than cis. Making use of the relation ∆G ) -RT ln Keq, the pKa for trans should be about 1.4 pH units higher than that for cis, which is close to the difference between two apparent pKa’s observed at 6.8 and 8. The formation of a hydroxyl radical-like oxidant from peroxynitrite depends upon peroxynitrite becoming protonated, isomerizing to the trans conformation, capturing enough energy to become vibrationally activated, and then reacting with an organic molecule. As a consequence, the rate of reaction is the same as the rate of peroxynitrite decomposition. While peroxynitrite is capable of forming a powerful oxidant with the reactivity of hydroxyl radical, this particular oxidative mechanism is not selective and is equally likely to randomly attack any molecule in a cell. The activitated form of peroxynitrite suffers from the same problem of free hydroxyl radical of being too reactive to be highly toxic. What is peculiar about peroxynitrite is its rather slow reaction with most biological molecules. Peroxynitrite is selective in what biological molecules it attacks. The half-life of peroxynitrite decomposition in phosphate buffer is about a second, which would enable peroxynitrite to undergo billions of collisions with cellular components without reaction. However, a few chemical groups do react directly with peroxynitrite, which favors selective reactions with key moieties in proteins, such as sulfhydryls, iron/sulfur centers, and zinc fingers. Peroxynitrite directly attacks sulfhydryls with second order rate constants that range from 1 × 103 to 6 × 103 M-1‚s-1 at 37 °C (11, 30). The pKa for the reaction with cysteine is at 6.8, consistent with a direct reaction between the cis conformation of peroxynitrite and thiols. While glutathione is an important thiol-containing antioxidant, the rate constants of 103 M-1‚s-1 imply that only one in a million collisions of peroxynitrite with the thiol on glutathione results in reaction. Physiological concentrations of glutathione will scavenge some peroxynitrite, but the reaction is not fast enough to completely protect other biological targets. When peroxynitrite is added to plasma, oxidation of proteins occurs simultaneously as endogenous antioxidants are being depleted (32). With most oxidants, the pool of antioxidants must be largely depleted before protein damage becomes apparent.
Beckman
Peroxynitrite can attack the same molecule by two different mechanisms. Pryor et al. have shown that peroxynitrite is a moderately good two electron oxidant of the amino acid methionine to form methionine sulfoxide, reacting by a second order reaction of 103 M-1‚s-1 (33). When the concentration of methionine drops below 1 mM, the principal product from methionine becomes ethylene. The formation of ethylene from methionine has long been used as an indicator of hydroxyl radical and is strong evidence for a one electron oxidation by peroxynitrite. When the concentration of methionine is not sufficiently high to react fast enough with peroxynitrite directly, a significant fraction of peroxynitrite can decompose through the hydroxyl radical-like pathway via free radical pathways. Peroxynitrite reacts particularly rapidly with ironsulfur centers and zinc fingers, with rate constants in excess of 105 M-1‚s-1 (34-36). The positive charge of the metal centers apparently helps to accelerate the oxidation of the sulfurs. These are among the most rapid reactions observed for peroxynitrite and are likely to be major targets in vivo. Indeed, peroxynitrite is a potent toxin to mitochondria (37, 38), which are rich in iron-sulfur centers involved in electron transfer. The pattern of inhibiting respiration is consistent with the inhibition of tumor cell respiration caused by activated macrophages (37, 39). Peroxynitrite can also trigger calcium release from mitochondria through a cyclosporin-inhibitable pathway (38). Transition metals like copper and iron also react directly with peroxynitrite and catalyze the nitration of tyrosines (21, 40). While such metals are often thought to be important for catalyzing free radical reactions, copper or iron catalyze a two electron oxidation in the reaction with peroxynitrite. The metal ion is more electropositive than hydrogen and pulls more electron density toward the metal, favoring heterolytic cleavage of the O-O bond.
Mn+ + -OONO f Mn-O‚‚‚ONO+ Thus, transition metals generate a species with a reactivity similar to nitronium ion (NO2+), a species generally produced at pH
