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We obviously agree with Wu et al. that H 2 O 2 might accumulate in the Archean land waters devoid of Fe 2+ . We do disagree on the topic of the half-life of H 2 O 2 , as the work cited in support for a longer half-live is not relevant to the conditions in the Archean ocean. While the existence of radicals in quartz is not in doubt, we do question the hypothesis that these radicals oxidize water to HO • and H 2 O 2 ., Competing Interests: Declaration of competing interest There is no conflict of interest., (Copyright © 2024 The Authors. Published by Elsevier B.V. All rights reserved.)

We address the chemical/biological history of H 2 O 2 back at the times of the Archean eon (2.5-3.9 billion years ago (Gya)). During the Archean eon the pO 2 was million-fold lower than the present pO 2 , starting to increase gradually from 2.3 until 0.6 Gya, when it reached ca. 0.2 bar. The observation that some anaerobic organisms can defend themselves against O 2 has led to the view that early organisms could do the same before oxygenic photosynthesis had developed at about 3 Gya. This would require the anaerobic generation of H 2 O 2 , and here we examine the various mechanisms which were suggested in the literature for this. Given the concentration of Fe 2+ at 20-200 μM in the Archean ocean, the estimated half-life of H 2 O 2 is ca. 0.7 s. The oceanic H 2 O 2 concentration was practically zero. We conclude that early organisms were not exposed to H 2 O 2 before the arrival of oxygenic photosynthesis., Competing Interests: Declaration of competing interest There is no conflict of interest., (Copyright © 2023 The Authors. Published by Elsevier B.V. All rights reserved.)

According to the literature, the Fenton reaction yields HO˙ and proceeds with 53 M -1 s -1 at 25 °C and low pH. Above pH 5, the reaction becomes first-order in HO - , and oxygen atom transfer has been detected, which indicates formation of oxidoiron(2+), FeO 2+ . These observations, and the assumption that the intermediate [FeHOO] + decays approximately iso-energetically to FeO 2+ , allow one to estimate an Gibbs energy of formation FeO 2+ of +15 ± 10 kJ mol -1 , from which follows the one-electron E °'(FeO 2+ , H 2 O/[Fe(HO) 2 ] + ) = +2.5 ± 0.1 V and the two-electron E °'(FeO 2+ , 2H + /Fe 2+ , H 2 O) = +1.36 ± 0.05 V, both at pH 7. In the presence of HCO 3 - , formation of FeCO 3 (aq) occurs which may facilitate formation of the [FeHOO] + intermediate, and leads to CO 3 ˙ - . At pH 7, the product of the Fenton reaction is thus FeO 2+ , or CO 3 ˙ - if HCO 3 - is present.

In an ab initio study, Tsuneda and Taketsugu ( Phys. Chem. Chem. Phys. , 2018, 20 , 24992-24999) discuss the Fenton reaction, the reaction of Fe 2+ with H 2 O 2 . They claim that reaction is endergonic and therefore introduce a new intermediate, a de facto monovalent iron complex. I show here that kinetically and thermodynamically such a monovalent iron complex cannot be formed.

In biology, the reaction of ONOO - with CO 2 is the main sink for ONOO - . This reaction yields CO 3 •- , NO 2 • , NO 3 - , and CO 2 . There is a long-standing debate with respect to the yield of the radicals relative to ONOO - . The reaction of ONOO - with CO 2 results at first in ONOOCO 2 - . According to one hypothesis, ONOOCO 2 - is extremely short-lived and devolves into a solvent cage that contains CO 3 •- and NO 2 • . Of these solvent cages, approximately two/thirds result in NO 3 - and CO 2 , and approximately one/third release CO 3 •- and NO 2 • that oxidize the substrate. According to our hypothesis, ONOOCO 2 - is formed much faster, is relatively long-lived, and may also be an oxidant; the limited yield is the result of ONOOCO 2 - being scavenged by a second CO 2 under conditions of a high CO 2 concentration. We disagree with the first hypothesis for three reasons: First, it is based on an estimated K for the reaction of ONOO - with CO 2 to form ONOOCO 2 - of ∼1 M -1 , while experiments yield a value of 4.5 × 10 3 M -1 . Second, we argue that the solvent cage as proposed is physically not realistic. Given the less than diffusion-controlled rate constant of CO 3 •- with NO 2 • , all radicals would escape from the solvent cage. Third, the reported ∼33% radical is not supported by an experiment where mass balance was established. We propose here a hybrid mechanism. After formation of ONOOCO 2 - , it undergoes homolysis to yield CO 3 •- with NO 2 • , or, depending on [CO 2 ], it is scavenged by a second CO 2 ; CO 3 •- oxidizes ONOO - , if present. These reactions allow us to successfully simulate the reaction of ONOO - with CO 2 over a wide range of ONOO - /CO 2 ratios. At lower ratios, fewer radicals are formed, while at higher ratios, radical yields between 30% and 40% are predicted. The differences in radical yields reported may thus be traced to the experimental ONOO - /CO 2 ratios. Given a physiological [CO 2 ] of 1.3 mM, the yield of CO 3 •- and NO 2 • is 19%, and lower if ONOOCO 2 - has a significant reactivity of its own.

A major form of toxicity arises from the ability of iron to redox cycle, that is, to accept an electron from a reducing compound and to pass it on to H 2 O 2 (the Fenton reaction). In order to do so, iron must be suitably complexed to avoid formation of Fe 2 O 3 . The ligands determine the electrode potential; this information should be known before experiments are carried out. Only one-electron transfer reactions are likely to be significant; thus two-electron potentials should not be used to determine whether an iron(III) complex can be reduced or oxidized. Ascorbate is the relevant reducing agent in blood serum, which means that iron toxicity in this compartment arises from the ascorbate-driven Fenton reaction. In the cytosol, an iron(II)-glutathione complex is likely to be the low-molecular weight iron complex involved in toxicity. When physiologically relevant concentrations are used the window of redox opportunity ranges from +0.1 V to +0.9 V. The electrode potential for non-transferrin-bound iron in the form of iron citrate is close to 0 V and the reduction of iron(III) citrate by ascorbate is slow. The clinically utilised chelators desferrioxamine, deferiprone and deferasirox in each case render iron complexes with large negative electrode potentials, thus being effective in preventing iron redox cycling and the associated toxicity resulting from such activity. There is still uncertainty about the product of the Fenton reaction, HO • or FeO 2+ ., (Copyright © 2018 Elsevier Inc. All rights reserved.)

Since 1981, Gordon Research Conferences have been held on the topic of Oxygen Radicals on a biennial basis, to highlight and discuss the latest cutting edge research in this area. Since the first meeting, one special feature of this conference has been the awarding of the so-called Iron Bolt, an award that started in jest but has gained increasing reputation over the years. Since no written documentation exists for this Iron Bolt award, this perspective serves to overview the history of this unusual award, and highlights various experiences of previous winners of this "prestigious" award and other interesting anecdotes., (Copyright © 2018 Elsevier Inc. All rights reserved.)

With CO 2 present in excess, ONOO - reacts to form an adduct in solution and in the solid state, most likely ONOOCO 2 - . In solution, the adduct appears within 2 ms and absorbs at 300 with an extinction coefficient, which is either 50% or 100% (preferred) of that of ONOO - , 1.70 × 10 3 M -1 cm -1 , and at 685 nm with an extinction coefficient of 85 M -1 cm -1 . When solid [(CH 3 ) 4 N][ONOO] is treated with CO 2 , these two maxima are red-shifted by 30-50 nm. The equilibrium constant for adduct formation in solution is (4.5 ± 0.5) × 10 3 M. The adduct reacts further with another CO 2 at a rate of (2.6 ± 0.8) × 10 4 M -1 s -1 and produces 2 CO 2 and NO 3 - . Thermochemical calculations show that ΟΝΟΟCO 2 - is a strong two-electron oxidizing agent, E°(ONOOCO 2 - , H + /NO 2 - , HCO 3 - ) = +1.28 V at pH 7 and an even stronger one-electron oxidizing agent E°'(ONOOCO 2 - , H + /NO 2 • , HCO 3 - ) = +1.51 V at pH 7. The extent of homolysis, that is formation of NO 2 • and CO 3 •- , is small, slightly less than 1% relative to ONOO - at the physiological concentration of CO 2 of 1.3 mM in plasma. Thus, ONOOCO 2 - is more relevant than CO 3 •- under in vivo conditions.

Cytochrome P450cam (CYP101Fe 3+ ) regioselectively hydroxylates camphor. Possible hydroxylating intermediates in the catalytic cycle of this well-characterized enzyme have been proposed on the basis of experiments carried out at very low temperatures and shunt reactions, but their presence has not yet been validated at temperatures above 0 °C during a normal catalytic cycle. Here, we demonstrate that it is possible to mimic the natural catalytic cycle of CYP101Fe 3+ by using pulse radiolysis to rapidly supply the second electron of the catalytic cycle to camphor-bound CYP101[FeO 2 ] 2+ Judging by the appearance of an absorbance maximum at 440 nm, we conclude that CYP101[FeOOH] 2+ (compound 0) accumulates within 5 μs and decays rapidly to CYP101Fe 3+ , with a k 440 nm of 9.6 × 10 4 s -1 All processes are complete within 40 μs at 4 °C. Importantly, no transient absorbance bands could be assigned to CYP101[FeO 2+ por •+ ] (compound 1) or CYP101[FeO 2+ ] (compound 2). However, indirect evidence for the involvement of compound 1 was obtained from the kinetics of formation and decay of a tyrosyl radical. 5-Hydroxycamphor was formed quantitatively, and the catalytic activity of the enzyme was not impaired by exposure to radiation during the pulse radiolysis experiment. The rapid decay of compound 0 enabled calculation of the limits for the Gibbs activation energies for the conversions of compound 0 → compound 1 → compound 2 → CYP101Fe 3+ , yielding a Δ G ‡ of 45, 39, and 39 kJ/mol, respectively. At 37 °C, the steps from compound 0 to the iron(III) state would take only 4 μs. Our kinetics studies at 4 °C complement the canonical mechanism by adding the dimension of time., (© 2017 by The American Society for Biochemistry and Molecular Biology, Inc.)

The autoxidation of NO • was studied in glass-like matrices of 2-methylbutane at 110 K and in a 8:3 v/v mixture of 2,2-dimethylbutane and n-pentane (rigisolve) at 80-90 K, by letting gaseous NO • diffuse into these solvents that were saturated with O 2 . In 2-methyllbutane, we observed a red compound. However, in rigisolve at 85-90 K, a bright yellow color appears that turns red when the sample is warmed by 10-20 K. The new yellow compound is a precursor of the red one and also diamagnetic. The UV-vis spectrum of the yellow compound contains a band which resembles that present in ONOO - . Because the red and yellow intermediates are not paramagnetic, we postulate that O═N-O-O • is in close contact with NO • , or with another O═N-O-O • . Diffusion of gaseous O 2 into rigisolve saturated with NO • does not produce a color; however, a weak EPR signal (g = 2.010) is observed. This signal most likely indicates the presence of ONOO • . These findings complement our earlier observation of a red color at low temperatures and the presence of ONOO • in the gas phase (Galliker, B.; Kissner, R.; Nauser, T.; Koppenol, W. H. Chem. Eur. J. 2009, 15, 6161-6168), and they indicate that the termolecular autoxidation of nitrogen monoxide proceeds via the intermediate ONOO • and not via N 2 O 2 .

There is currently interest in sulfur-containing molecules that may or may not play a role in signaling. We have collected relevant thermodynamic data, namely standard Gibbs energies of formation and electrode potentials at pH 7, and used these to construct a Frost diagram. Thermodynamic data not available in the literature could be estimated with reasonable confidence. At pH 7, the electrode potential of the RSS/RSS - couple is +0.68 V, 0.28 V less than that of the RS, H + /RSH couple. S 2 - is unstable with respect to HSS - and S 2 . Generally, polysulfur compounds, with the exception of RSSR, are thermodynamically unstable with respect to disproportionation and ultimately lead to formation of kinetically inert S 8(s) . About thermoneutral is the formation of RSS - from RSSR and HS - , but formation of HNO from HS - and SNO - , and from HS - and RSNO, is unfavorable. The formation of SSNO - in vivo is kinetically unlikely., (Copyright © 2016 Elsevier Inc. All rights reserved.)

Electrode potentials for aromatic amino acid radical/amino acid couples were deduced from cyclic voltammograms and pulse radiolysis experiments. The amino acids investigated were l-tryptophan, l-tyrosine, N-acetyl-l-tyrosine methyl ester, N-acetyl-3-nitro-l-tyrosine ethyl ester, N-acetyl-2,3-difluoro-l-tyrosine methyl ester, and N-acetyl-2,3,5-trifluoro-l-tyrosine methyl ester. Conditional potentials were determined at pH 7.4 for all compounds listed; furthermore, Pourbaix diagrams for l-tryptophan, l-tyrosine, and N-acetyl-3-nitro-l-tyrosine ethyl ester were obtained. Electron transfer accompanied by proton transfer is reversible, as confirmed by detailed analysis of the current waves, and because the slopes of the Pourbaix diagrams obey Nernst's law. E°'(Trp(•),H(+)/TrpH) and E°'(TyrO(•),H(+)/TyrOH) at pH 7 are 0.99 ± 0.01 and 0.97 ± 0.01 V, respectively. Pulse radiolysis studies of two dipeptides that contain both amino acids indicate a difference in E°' of approximately 0.06 V. Thus, in small peptides, we recommend values of 1.00 and 0.96 V for E°'(Trp(•),H(+)/TrpH) and E°'(TyrO(•),H(+)/TyrOH), respectively. The electrode potential of N-acetyl-3-nitro-l-tyrosine ethyl ester is higher, while because of mesomeric stabilization of the radical, those of N-acetyl-2,3-difluoro-l-tyrosine methyl ester and N-acetyl-2,3,5-trifluoro-l-tyrosine methyl ester are lower than that of tyrosine. Given that the electrode potentials at pH 7 of E°'(Trp(•),H(+)/TrpH) and E°'(TyrO(•),H(+)/TyrOH) are nearly equal, they would be, in principle, interchangeable. Proton-coupled electron transfer pathways in proteins that use TrpH and TyrOH are thus nearly thermoneutral.

A history of the formation of hydrogen peroxide in vivo is presented, starting with the discovery of catalase. The first hypothesis was formulated by Heinrich Wieland, who assumed that dioxygen reacted directly with organic molecules. This view was strongly criticised by Otto Warburg, Helmut Sies' academic grandfather. The involvement of hydrogen peroxide in physiological processes was investigated by Theodor Bücher, the "Doktorvater" of Helmut. Helmut's research made it possible to quantitate hydrogen peroxide in tissues., (Copyright © 2015 Elsevier Inc. All rights reserved.)

Iron in iron overload disease is present as non-transferrin-bound iron, consisting of iron, citrate, and albumin. We investigated the redox properties of iron citrate by electrochemistry, by the kinetics of its reaction with ascorbate, by ESR, and by analyzing the products of reactions of ascorbate with iron citrate complexes in the presence of H2O2 with 4-hydroxybenzoic acid as a reporter molecule for hydroxylation. We report -0.03 V < E°' > +0.01 V for the (Fe(3+)-cit/Fe(2+)-cit) couple. The first step in the reaction of iron citrate with ascorbate is the rapid formation of mixed complexes of iron with citrate and ascorbate, followed by slow reduction to Fe(2+)-citrate with k = ca. 3 M(-1) s(-1). The ascorbyl radical is formed by iron citrate oxidation of Hasc(-) with k = ca. 0.02 M(-1) s(-1); the majority of the ascorbyl radical formed is sequestered by complexation with iron and remains EPR silent. The hydroxylation of 4-hydroxybenzoic acid driven by the Fenton reduction of iron citrate by ascorbate in the presence of H2O2 proceeds in three phases: the first phase, which is independent of the presence of O2, is revealed as a nonzero intercept that reflects the rapid reaction of accumulated Fe(2+) with H2O2; the intermediate oxygen-dependent phase fits a first-order accumulation of product with k = 5 M(-1) s(-1) under aerobic and k = 13 M(-1) s(-1) under anaerobic conditions; the slope of the final linear phase is ca. k = 5 × 10(-2) M(-1) s(-1) under both aerobic and anaerobic conditions. Product yields under aerobic conditions are greater than predicted from the initial concentration of iron, but they are less than predicted for continuous redox cycling in the presence of excess ascorbate. The ongoing formation of hydroxylated product supports slow redox cycling by iron citrate. Thus, when H2O2 is available, iron-citrate complexes may contribute to pathophysiological manifestations of iron overload diseases.

Protein thiyl radicals are important intermediates generated in redox processes of thiols and disulfides. Thiyl radicals efficiently react with glutathione and ascorbate, and the common notion is that these reactions serve to eliminate thiyl radicals before they can enter potentially hazardous processes. However, over the past years increasing evidence has been provided for rather efficient intramolecular hydrogen transfer processes of thiyl radicals in proteins and peptides. Based on rate constants published for these processes, we have performed kinetic simulations of protein thiyl radical reactivity. Our simulations suggest that protein thiyl radicals enter intramolecular hydrogen transfer reactions to a significant extent even under physiologic conditions, i.e., in the presence of 30 µM oxygen, 1 mM ascorbate, and 10 mM glutathione. At lower concentrations of ascorbate and glutathione, frequently observed when tissue is exposed to oxidative stress, the extent of irreversible protein thiyl radical-dependent protein modification increases., (Copyright © 2014 Elsevier Inc. All rights reserved.)

Cytochrome P450 3A4 (CYP3A4) is the major human P450 responsible for the metabolism of carbamazepine (CBZ). To explore the mechanisms of interactions of CYP3A4 with this anticonvulsive drug, we carried out multiple molecular dynamics (MD) simulations, starting with the complex of CYP3A4 manually docked with CBZ. On the basis of these simulations, we engineered CYP3A4 mutants I369F, I369L, A370V, and A370L, in which the productive binding orientation was expected to be stabilized, thus leading to increased turnover of CBZ to the 10,11-epoxide product. In addition, we generated CYP3A4 mutant S119A as a control construct with putative destabilization of the productive binding pose. Evaluation of the kinetics profiles of CBZ epoxidation demonstrate that CYP3A4-containing bacterial membranes (bactosomes) as well as purified CYP3A4 (wild-type and mutants I369L/F) exhibit substrate inhibition in reconstituted systems. In contrast, mutants S119A and A370V/L exhibit S-shaped profiles that are indicative of homotropic cooperativity. MD simulations with two to four CBZ molecules provide evidence that the substrate-binding pocket of CYP3A4 can accommodate more than one molecule of CBZ. Analysis of the kinetics profiles of CBZ metabolism with a model that combines the formalism of the Hill equation with an allowance for substrate inhibition demonstrates that the mechanism of interactions of CBZ with CYP3A4 involves multiple substrate-binding events (most likely three). Despite the retention of the multisite binding mechanism in the mutants, functional manifestations reveal an exquisite sensitivity to even minor structural changes in the binding pocket that are introduced by conservative substitutions such as I369F, I369L, and A370V.

Beta-glucan is a dietary fiber, which possesses several health benefits, such as cholesterol lowering, however this fiber is easily degraded in the presence of Fenton reagents. In the present study, the iron binding capacity of oat beta-glucan and barley beta-glucan was evaluated by investigating the kinetics of the Fenton reaction at pH 2.7 and 4.7 using stopped flow spectroscopy. The rate constant of the Fenton reaction is not affected by the presence of the beta-glucans in a solution pH 2.7, hence none of the beta-glucans bind iron(II) at this pH. However, at pH 4.7, the kinetics of the Fenton reaction vary between acetate buffer (k=2.8×10(2)M(-1)s(-1)), barley beta-glucan (k=2.2×10(2)M(-1)s(-1)) and oat beta-glucan (k=1.2×10(2)M(-1)s(-1)), which demonstrates that barley beta-glucan and oat beta-glucan form complexes with iron(II). Moreover, oat beta-glucan shows a stronger affinity for iron(II) than barley beta-glucan, and may thereby reduce the formation of hydroxyl radicals and diminish the rate of viscosity loss of the oat beta-glucan solution, as shown by the ESR and rheological data. The results presented in this study suggest that cereal beta-glucans can potentially reduce the bioavailability of iron., (Copyright © 2014 Elsevier Ltd. All rights reserved.)

H2 has been suggested to act as an antioxidant when administered just before the reperfusion phase of induced oxidative stress. These effects have been reported, for example, for the heart, brain, and liver. It is hypothesized that this beneficial effect may be due to selective scavenging of HO(⋅) and ONOOH by H2. The reaction of H2 with HO(⋅) has been studied by pulse radiolysis in the past and is too slow to be physiologically relevant, not to mention that the reaction yields the reactive H(⋅) radical. We therefore investigated whether H2 reacts with ONOOH and whether the presence of H2 influences the yield of nitration of tyrosine by ONOOH. With only negative results, we entertained the notion that H2 may possibly exert its beneficial effects by reducing Fe(III) centers, oxidized during oxidative stress. However, neither hemes nor iron-sulfur clusters were reduced., (Copyright © 2014 Elsevier Inc. All rights reserved.)

Thioredoxin reductases, important biological redox mediators for two-electron transfers, contain either 2 cysteines or a cysteine (Cys) and a selenocysteine (Sec) at the active site. The incorporation of Sec is metabolically costly, and therefore surprising. We provide here a rationale: in the case of an accidental one-electron transfer to a S-S or a S-Se bond during catalysis, a thiyl or a selanyl radical, respectively would be formed. The thiyl radical can abstract a hydrogen from the protein backbone, which subsequently leads to the inactivation of the protein. In contrast, a selanyl radical will not abstract a hydrogen. Therefore, formation of Sec radicals in a GlyCysSecGly active site will less likely result in the destruction of a protein compared to a GlyCysCysGly active site.

Tyrosine (Tyr) residues are major sites of radical generation during protein oxidation. We used insulin as a model to study the kinetics, mechanisms, and products of the reactions of radiation-induced or enzyme-generated protein-tyrosyl radicals with superoxide to demonstrate the feasibility of these reactions under oxidative stress conditions. We found that insulin-tyrosyl radicals combined to form dimers, mostly via the tyrosine at position 14 on the α chain (Tyr14). However, in the presence of superoxide, dimerization was largely outcompeted by the reaction of superoxide with insulin-tyrosyl radicals. Using pulse radiolysis, we measured a second-order rate constant for the latter reaction of (6±1) × 10(8) M(-1) s(-1) at pH 7.3, representing the first measured rate constant for a protein-tyrosyl radical with superoxide. Mass-spectrometry-based product analyses revealed the addition of superoxide to the insulin-Tyr14 radical to form the hydroperoxide. Glutathione efficiently reduced the hydroperoxide to the corresponding monoxide and also subsequently underwent Michael addition to the monoxide to give a diglutathionylated protein adduct. Although much slower, conjugation of the backbone amide group can form a bicyclic Tyr-monoxide derivative, allowing the addition of only one glutathione molecule. These findings suggest that Tyr-hydroperoxides should readily form on proteins under oxidative stress conditions where protein radicals and superoxide are both generated and that these should form addition products with thiol compounds such as glutathione., (Copyright © 2014 Elsevier Inc. All rights reserved.)

The reactions of NO2 with both oxidized and reduced cytochrome c at pH 7.2 and 7.4, respectively, and with N-acetyltyrosine amide and N-acetyltryptophan amide at pH 7.3 were studied by pulse radiolysis at 23 °C. NO2 oxidizes N-acetyltyrosine amide and N-acetyltryptophan amide with rate constants of (3.1±0.3)×10(5) and (1.1±0.1)×10(6) M(-1) s(-1), respectively. With iron(III)cytochrome c, the reaction involves only its amino acids, because no changes in the visible spectrum of cytochrome c are observed. The second-order rate constant is (5.8±0.7)×10(6) M(-1) s(-1) at pH 7.2. NO2 oxidizes iron(II)cytochrome c with a second-order rate constant of (6.6±0.5)×10(7) M(-1) s(-1) at pH 7.4; formation of iron(III)cytochrome c is quantitative. Based on these rate constants, we propose that the reaction with iron(II)cytochrome c proceeds via a mechanism in which 90% of NO2 oxidizes the iron center directly-most probably via reaction at the solvent-accessible heme edge-whereas 10% oxidizes the amino acid residues to the corresponding radicals, which, in turn, oxidize iron(II). Iron(II)cytochrome c is also oxidized by peroxynitrite in the presence of CO2 to iron(III)cytochrome c, with a yield of ~60% relative to peroxynitrite. Our results indicate that, in vivo, NO2 will attack preferentially the reduced form of cytochrome c; protein damage is expected to be marginal, the consequence of formation of amino acid radicals on iron(III)cytochrome c., (Copyright © 2014 Elsevier Inc. All rights reserved.)

Production of minute concentrations of superoxide (O2(*-)) and nitrogen monoxide (nitric oxide, NO*) plays important roles in several aspects of cellular signaling and metabolic regulation. However, in an inflammatory environment, the concentrations of these radicals can drastically increase and the antioxidant defenses may become overwhelmed. Thus, biological damage may occur owing to redox imbalance-a condition called oxidative and/or nitrosative stress. A complex interplay exists between iron metabolism, O2(*-), hydrogen peroxide (H2O2), and NO*. Iron is involved in both the formation and the scavenging of these species. Iron deficiency (anemia) (ID(A)) is associated with oxidative stress, but its role in the induction of nitrosative stress is largely unclear. Moreover, oral as well as intravenous (iv) iron preparations used for the treatment of ID(A) may also induce oxidative and/or nitrosative stress. Oral administration of ferrous salts may lead to high transferrin saturation levels and, thus, formation of non-transferrin-bound iron, a potentially toxic form of iron with a propensity to induce oxidative stress. One of the factors that determine the likelihood of oxidative and nitrosative stress induced upon administration of an iv iron complex is the amount of labile (or weakly-bound) iron present in the complex. Stable dextran-based iron complexes used for iv therapy, although they contain only negligible amounts of labile iron, can induce oxidative and/or nitrosative stress through so far unknown mechanisms. In this review, after summarizing the main features of iron metabolism and its complex interplay with O2(*-), H2O2, NO*, and other more reactive compounds derived from these species, the potential of various iron therapies to induce oxidative and nitrosative stress is discussed and possible underlying mechanisms are proposed. Understanding the mechanisms, by which various iron formulations may induce oxidative and nitrosative stress, will help us develop better tolerated and more efficient therapies for various dysfunctions of iron metabolism., (© 2013 Elsevier Inc. All rights reserved.)

Wegerich et al. (J. Biol. Inorg. Chem. 18:429-440, 2013), working with singly modified human cytochromes c, claim to have found a new mechanism for the reduction of iron(III) cytochrome c by superoxide. I show that electron transfer by way of the solvent-accessible haem edge-a mechanism not considered by Wegerich et al.-is still the correct mechanism. Furthermore, several deficiencies in this work preclude any comparisons with other publications on this topic.

The decay of ONOOH near neutral pH has been examined as a function of isomerization to H(+) and NO3(-), and decomposition to NO2(-) and O2via O2NOO(-). We find that in phosphate buffer k(isomerization) = 1.11 ± 0.01 s(-1) and k(disproportionation) = (1.3 ± 0.1) × 10(3) M(-1) s(-1) at 25 °C and I = 0.2 M. In the presence of 0.1 M tris(hydroxymethyl)aminomethane (Tris), the decay proceeds more rapidly: k(disproportionation) = 9 × 10(3) M(-1) s(-1). The measured first half-life of the absorbance of peroxynitrite correlates with [Tris]0·([ONOO(-)]0 + [ONOOH]0)(2), where the subscript 0 indicates initial concentrations; if this function exceeds 6.3 × 10(-12) M(3), then Tris significantly accelerates the decomposition of peroxynitrite.

Poly(α-methylstyrene sulfonic acid) (PAMS) represents a class of polymers that can form the protogenic constituent in electrolyte membranes for fuel cells. Oxidative stress is thought to play an important role in the degradation of the fuel cell membranes. Having previously established that damage may be mediated via abstraction of a benzylic hydrogen, we examined model compounds similar to those used before, but with a methyl group at the α-position. We studied the reaction of HO˙ and SO4(˙-), generated by pulse radiolysis, with model compounds in aqueous solution, and measured k = (2 ± 0.5) × 10(10) M(-1) s(-1) and (2 - 3) × 10(10) M(-1) s(-1) for the reaction of HO˙ with PAMS with average molecular weights of 2640 Da (PAMS-2640) and 6440 Da (PAMS-6440), respectively, at room temperature. At low pH, the decay of the hydroxycyclohexadienyl radical thus formed is accompanied by the formation of an absorption band in the visible region of the spectrum, which we tentatively assign to the radical cation of PAMS-2640 and -6440. The radical cation of PAMS-2640, formed by the reaction of SO4(˙-) with k = (6 ± 1) × 10(8) M(-1) s(-1), has a local absorption maximum at 560 nm, with ε560 ≥ 1400 M(-1) cm(-1). For the reaction of HO˙ and SO4(˙-) with the model compound benzenesulfonate, we measured k = (4-5) × 10(9) M(-1) s(-1) and (1.0 ± 0.3) × 10(8) M(-1) s(-1), respectively, while the reaction of SO4(˙-) with PAMS-6440 proceeds with (0.8-1) × 10(9) M(-1) s(-1). The 4-sulfophenoxyl radical was generated via the reaction of N3˙ with 4-hydroxybenzenesulfonate; ε410 ≥ 2300 M(-1) cm(-1). Not unexpectedly, the radical cation of PAMS is longer-lived than that of polystyrene sulfonic acid. Furthermore, fragmentation may result in desulfonation.

The isomerisation of ONOOH to NO(3)(-) and H(+), some oxidations and all hydroxylations and nitrations of aromatic compounds are first-order in ONOOH and zero-order in the compounds that are modified. These reactions are widely believed to proceed via homolysis of ONOOH into HO˙ and NO(2)˙ to an extent of ca. 30%. We review the evidence pro and contra homolysis in studies that involve (1) thermochemical considerations, (2) isomerisation to NO(3)(-) and H(+), (3) decomposition to NO(2)(-) and O(2), (4) HO˙ scavenger studies, (5) deuterium isotope effects, (6) (18)O-scrambling studies, (7) electrochemistry, (8) CIDNP NMR, and (9) photolysis. Our conclusion is that homolysis may be involved to a minor extent of ca. 5%. The initiation of ONOOH isomerisation may be visualised as an out-of-plane vibration of the terminal HO-group relative to the nitrogen. At ONOO(-) concentrations exceeding 0.1 mM and near neutral pH, disproportionation to NO(2)(-) and O(2) occurs; such disproportionations are typical for peroxy acids. For oxidation and nitration of organic substrates, we favour a mechanism involving initial formation of an adduct between the compound to be oxidised or nitrated and ONOOH.

Ascorbate levels decrease in organisms subjected to oxidative stress, but the responsible reactions have not been identified. Our earlier studies have shown that protein C-centered radicals react rapidly with ascorbate. In aerobes, these radicals can react with oxygen to form peroxyl radicals. To estimate the relative probabilities of the reactions of ascorbate with protein C- and O-centered radicals, we measured by pulse radiolysis the rate constants of the reactions of C-centered radicals in Gly, Ala, and Pro with O₂ and of the resultant peroxyl radicals with ascorbate. Calculations based on the concentrations of ascorbate and oxygen in human tissues show that the relative probabilities of reactions of the C-centered amino acid radicals with O₂ and ascorbate vary between 1:2.6 for the pituitary gland and 1:0.02 for plasma, with intermediate ratios for other tissues. The high frequency of occurrence of Gly, Ala, and Pro in proteins and the similar reaction rate constants of their C-centered radicals with O₂ and their peroxo-radicals with ascorbate suggest that our results are also valid for proteins. Thus, the formation of protein C- or O-centered radicals in vivo can account for the loss of ascorbate in organisms under oxidative stress., (Copyright © 2012 Elsevier Inc. All rights reserved.)

Cysteine, cysteinyl-glycine, glutathione, phenylalanyl-cysteinyl-glycine, and histidyl-cysteinyl-glycine were dissolved in acidic and neutral D(2)O in the presence of the radical generator 2,2'-azobis(2-methylpropionamidine) dihydrochloride and radical mediator compounds (benzyl alcohol and 2-propanol). An exchange of H-atoms by D-atoms took place in these peptides due to intramolecular H-abstraction equilibria. NMR measurements allow one to follow the extent of H-D exchanges and to identify the sites where these exchanges take place. Significant exchanges occur in acidic media in GSH at positions Glu-β and Glu-γ, in Phe-Cys-Gly at positions Phe ortho, Phe-β, Cys-α, Cys-β, and Gly-α, and in His-Cys-Gly at positions His H1, His H2, His β, Cys β, and Gly α. In neutral media, exchanges occur in Cys-Gly at position Cys β and in GSH at position Cys α. Phe-Cys-Gly and His-Cys-Gly were not examined in neutral media. Sites participating in the radical exchange equilibria are highly dependent on structure and pH; the availability of electron density in the form of lone pairs appears to increase the extent of exchange. Interestingly, and unexpectedly, 2D NMR experiments show that GSH rearranges itself in acidic solution: the signals shift, but their patterns do not change. The formation of a thiolactone from Gly and Cys residues matches the changes observed.

Dihydrogen sulfide recently emerged as a biological signaling molecule with important physiological roles and significant pharmacological potential. Chemically plausible explanations for its mechanisms of action have remained elusive, however. Here, we report that H(2)S reacts with S-nitrosothiols to form thionitrous acid (HSNO), the smallest S-nitrosothiol. These results demonstrate that, at the cellular level, HSNO can be metabolized to afford NO(+), NO, and NO(-) species, all of which have distinct physiological consequences of their own. We further show that HSNO can freely diffuse through membranes, facilitating transnitrosation of proteins such as hemoglobin. The data presented in this study explain some of the physiological effects ascribed to H(2)S, but, more broadly, introduce a new signaling molecule, HSNO, and suggest that it may play a key role in cellular redox regulation.

Nitrosothiols are powerful vasodilators. Although the mechanism of their formation near neutral pH is an area of intense research, neither the energetics nor the kinetics of this reaction or of subsequent reactions have been addressed. The following considerations may help to guide experiments. (1) The standard Gibbs energy for the homolysis reaction RSNO → RS(•) + NO(•)(aq) is +110 ± 5 kJ mol(-1). (2) The electrode potential of the RSNO, H(+)/RSH, NO(•)(aq) couple is -0.20 ± 0.06 V at pH 7. (3) Thiol nitrosation by NO(2)(-) is favorable by 37 ± 5 kJ mol(-1) at pH 7. (4) N(2)O(3) is not involved in in vivo nitrosation mechanisms for thermodynamic--its formation from NO(2)(-) costs 59 kJ mol(-1)--or kinetic--the reaction being second-order in NO(2)(-)--reasons. (5) Hemoglobin (Hb) cannot catalyze formation of N(2)O(3), be it via the intermediacy of the reaction of Hb[FeNO(2)](2+) with NO(•) (+81 kJ mol(-1)) or reaction of Hb[FeNO](3+) with NO(2)(-) (+88 kJ mol(-1)). (6) Energetically and kinetically viable are nitrosations that involve HNO(2) or NO(•) in the presence of an electron acceptor with an electrode potential higher than -0.20 V. These considerations are derived from existing thermochemical and kinetics data.

The mercapto group of cysteine (Cys) is a predominant target for oxidative modification, where one-electron oxidation leads to the formation of Cys thiyl radicals, CysS(•). These Cys thiyl radicals enter 1,2- and 1,3-hydrogen transfer reactions, for which rate constants are reported in this paper. The products of these 1,2- and 1,3-hydrogen transfer reactions are carbon-centered radicals at position C(3) (α-mercaptoalkyl radicals) and C(2) ((•)C(α) radicals) of Cys, respectively. Both processes can be monitored separately in Cys analogues such as cysteamine (CyaSH) and penicillamine (PenSH). At acidic pH, thiyl radicals from CyaSH permit only the 1,2-hydrogen transfer according to equilibrium 12, (+)H(3)NCH(2)CH(2)S(• )⇌ (+)H(3)NCH(2)(•)CH-SH, where rate constants for forward and reverse reaction are k(12) ≈ 10(5) s(-1) and k(-12) ≈ 1.5 × 10(5)s(-1), respectively. In contrast, only the 1,3-hydrogen transfer is possible for thiyl radicals from PenSH according to equilibrium 14, ((+)H(3)N/CO(2)H)C(α)-C(CH(3))(2)-S(•) ⇌ ((+)H(3)N/CO(2)H)(•)C(α)-C(CH(3))(2)-SH, where rate constants for the forward and the reverse reaction are k(14) = 8 × 10(4) s(-1) and k(-14) = 1.4 × 10(6) s(-1). The (•)C(α) radicals from PenSH and Cys have the additional opportunity for β-elimination of HS(•)/S(•-), which proceeds with k(39) ≈ (3 ± 1) × 10(4) s(-1) from (•)C(α) radicals from PenSH and k(-34) ≈ 5 × 10(3) s(-1) from (•)C(α) radicals from Cys. The rate constants quantified for the 1,2- and 1,3-hydrogen transfer reactions can be used as a basis to calculate similar processes for Cys thiyl radicals in proteins, where hydrogen transfer reactions, followed by the addition of oxygen, may lead to the irreversible modification of target proteins.

The repair of tryptophan and tyrosine radicals in proteins by urate was studied by pulse radiolysis. In chymotrypsin, urate repairs tryptophan radicals efficiently with a rate constant of 2.7 × 10(8)M(-1)s(-1), ca. 14 times higher than the rate constant derived for N-acetyltryptophan amide, 1.9 × 10(7)M(-1)s(-1). In contrast, no repair of tryptophan radicals was observed in pepsin, which indicates a rate constant smaller than 6 × 10(7)M(-1)s(-1). Urate repairs tyrosine radicals in pepsin with a rate constant of 3 × 10(8)M(-1)s(-1)-ca. 12 times smaller than the rate constant reported for free tyrosine-but not in chymotrypsin, which implies an upper limit of 1 × 10(6)M(-1)s(-1) for the corresponding rate constant. Intra- and intermolecular electron transfer from tyrosine residues to tryptophan radicals is observed in both proteins, however, to different extents and with different rate constants. Urate inhibits electron transfer in chymotrypsin but not in pepsin. Our results suggest that urate repairs the first step on the long path to protein modification and prevents damage in vivo. It may prove to be a very important repair agent in tissue compartments where its concentration is higher than that of ascorbate. The product of such repair, the urate radical, can be reduced by ascorbate. Loss of ascorbate is then expected to be the net result, whereas urate is conserved., (Copyright © 2012 Elsevier Inc. All rights reserved.)

Selenocysteine is present in a variety of proteins and catalyzes the oxidation of thiols to disulfides and the reduction of disulfides to thiols. Here, we compare the kinetic and thermodynamic properties of cysteine with its selenium-containing analogon, selenocysteine. Reactions of simple selenols at pH 7 are up to four orders of magnitude faster than their sulfur analogs, depending on reaction type. In redox-related proteins, the use of selenium instead of sulfur can be used to tune electrode, or redox, potentials. Selenocysteine could also have a protective effect in proteins because its one-electron oxidized product, the selanyl radical, is not oxidizing enough to modify or destroy proteins, whereas the cysteine-thiyl radical can do this very rapidly.

Manganese(III)-meso-tetraphenylporphyrin [Mn(TPP)] and manganese(III)-meso-tetrakis(pentafluorophenyl)porphyrin [Mn(TPFPP)] catalyse the epoxidation of cyclooctene by IO(4)(-) in the presence of excess imidazoles, in both dry CH(2)Cl(2) and CH(2)Cl(2) saturated with H(2)O. The reaction rates of the electron deficient Mn(TPFPP) are a factor 24 less than those of Mn(TPP); however, the former increases 15-30 times in the presence of water, while those of Mn(TPP) do so by a factor of 2-3. The most striking catalytic enhancement caused by the addition of water was observed with 2-methylimidazole and Mn(TPFPP). As deprotonation of imidazoles may play a significant role in the presence of water, we found that manganese(III)-meso-tetrakis(phenyl-4-sulfonato)porphyrin [Mn(TPPS)] decreases the NH proton pK(a) of axially coordinated imidazole from 14.2 to 9.5. We conclude that the imidazole ligand is partially deprotonated in the presence of water. The latter enables the solvation of imidazolium ions that are formed simultaneously. The imidazolate form of the co-catalyst is a much stronger donor than the imidazole itself, providing electron density to Mn(III) and thus promoting oxygen transfer. The failure of N-methylimidazole to increase the reaction rates upon addition of water supports this hypothesis. A functionally related deprotonation has been shown to occur in horseradish peroxidase (J. S. de Ropp, V. Thanabal, G. N. La Mar, J. Am. Chem. Soc. 1985, 107, 8270-8272) and in chlorite dismutase (B. R. Goblirsch, B. R. Streit, J. L. Dubois, C. M. Wilmot, J. Biol. Inorg. Chem. 2010, 15, 879-888). Mn(III)porphyrins in combination with imidazoles and water constitute a functional biomimetic model of peroxidases., (This journal is © The Royal Society of Chemistry 2011)

Clarification of the mechanism of degradation of model compounds for polymers used in polymer electrolyte fuel cells may identify intermediates that propagate damage; such knowledge can be used to improve the lifetime of fuel cell membranes, a central issue to continued progress in fuel cell technology. In proton-exchange membranes based on poly(styrene sulfonic acid), hydroxycyclohexadienyl radicals are formed after reaction with HO˙ and thought to decay to short-lived radical cations at low pH. To clarify subsequent reactions, we generated radical cations by reaction of SO(4)˙(-) with oligomers of poly(styrene sulfonic acid) (MW ≈ 1100 Da). At 295 K, this reaction proceeds with k = (4.5 ± 0.6) × 10(8) M(-1) s(-1), both at pH 2.4 and 3.4, and yields benzyl radicals with an estimated yield of ≤60% relative to [SO(4)˙(-)]. The radical cation is too short-lived to be observed: based on a benzyl radical yield of 60%, a lower limit of k > 6.8 × 10(5) s(-1) for the intramolecular transformation of the aromatic radical cation of the oligomer to a benzyl radical is deduced. Our results show that formation of the benzyl radical, an important precursor in the breakdown of the polymer, is irreversible., (This journal is © the Owner Societies 2011)

Otto Warburg pioneered quantitative investigations of cancer cell metabolism, as well as photosynthesis and respiration. Warburg and co-workers showed in the 1920s that, under aerobic conditions, tumour tissues metabolize approximately tenfold more glucose to lactate in a given time than normal tissues, a phenomenon known as the Warburg effect. However, this increase in aerobic glycolysis in cancer cells is often erroneously thought to occur instead of mitochondrial respiration and has been misinterpreted as evidence for damage to respiration instead of damage to the regulation of glycolysis. In fact, many cancers exhibit the Warburg effect while retaining mitochondrial respiration. We re-examine Warburg's observations in relation to the current concepts of cancer metabolism as being intimately linked to alterations of mitochondrial DNA, oncogenes and tumour suppressors, and thus readily exploitable for cancer therapy.

The fast reaction of (•)NO and O(2)(•-) to give ONOO(-) has been extensively studied at irreversible conditions, but the reasons for the wide variations in observed forward rate constants (3.8 ≤ k(f) ≤ 20 × 10(9) M(-1) s(-1)) remain unexplained. We characterized the diffusion-dependent aqueous (pH > 12) chemical equilibrium of the form (•)NO + O(2)(•-) = ONOO(-) with respect to its dependence on temperature, viscosity, and [ONOO(-)](eq) by determining [ONOO(-)](eq) and [(•)NO](eq). The equilibrium forward reaction rate constant (k(f)(eq)) has negative activation energy, in contrast to that found under irreversible conditions. In contradiction to the law of mass action, we demonstrate that the equilibrium constant depends on ONOO(-) concentration. Therefore, a wide range of k(f)(eq) values could be derived (7.5-21 × 10(9) M(-1) s(-1)). Of general interest, the variations in k(f) can thus be explained by its dependence on the distance between ONOO(-) particles (sites of generation of (•)NO and O(2)(•-)).

The oxidation of proteins and other macromolecules by radical species under conditions of oxidative stress can be modulated by antioxidant compounds. Decreased levels of the antioxidants glutathione and ascorbate have been documented in oxidative stress-related diseases. A radical generated on the surface of a protein can: (1) be immediately and fully repaired by direct reaction with an antioxidant; (2) react with dioxygen to form the corresponding peroxyl radical; or (3) undergo intramolecular long range electron transfer to relocate the free electron to another amino acid residue. In pulse radiolysis studies, in vitro production of the initial radical on a protein is conveniently made at a tryptophan residue, and electron transfer often leads ultimately to residence of the unpaired electron on a tyrosine residue. We review here the kinetics data for reactions of the antioxidants glutathione, selenocysteine, and ascorbate with tryptophanyl and tyrosyl radicals as free amino acids in model compounds and proteins. Glutathione repairs a tryptophanyl radical in lysozyme with a rate constant of (1.05±0.05)×10(5) M(-1) s(-1), while ascorbate repairs tryptophanyl and tyrosyl radicals ca. 3 orders of magnitude faster. The in vitro reaction of glutathione with these radicals is too slow to prevent formation of peroxyl radicals, which become reduced by glutathione to hydroperoxides; the resulting glutathione thiyl radical is capable of further radical generation by hydrogen abstraction. Although physiologically not significant, selenoglutathione reduces tyrosyl radicals as fast as ascorbate. The reaction of protein radicals formed on insulin, β-lactoglobulin, pepsin, chymotrypsin and bovine serum albumin with ascorbate is relatively rapid, competes with the reaction with dioxygen, and the relatively innocuous ascorbyl radical is formed. On the basis of these kinetics data, we suggest that reductive repair of protein radicals may contribute to the well-documented depletion of ascorbate in living organisms subjected to oxidative stress.

The reduction of oxidized glutathione GSSG by hydrated electrons and hydrogen atoms to form GSSG•⁻ is quantitative. The radical anion dissociates into GS• and GS⁻, and the S-centered radical subsequently abstracts a hydrogen intramolecularly. We observe sequential development of UV absorbance signatures that indicate the formation of both α- and β-carbon-centered radicals. From experiments performed at pH 2 and pH 11.8, we determined forward and reverse rate constants for the overall equilibrium between sulfur-centered and carbon-centered radicals: k(forward) = 3·10⁵ s⁻¹, k(reverse) = 7·10⁵ s⁻¹, and K = 0.4. Furthermore, on the basis of the differences between the kinetics traces at 240 and 280 nm, we estimate that α- and β-carbon-centered radicals are formed at a surprising ratio of 1:3. The ratios found at pH 2 also apply to pH 7, with the conclusion that the equilibrium ratio of S-centered:β-centered:α-centered radicals is, very approximately, 8:3:1. The formation of carbon-centered radicals could lead to irreversible damage in proteins via the formation of carbon-carbon bonds or backbone fragmentation.

An understanding of the reactivity of oligomeric compounds that model fuel cell membrane materials under oxidative-stress conditions that mimic the fuel cell operating environment can identify material weaknesses and yield valuable insights into how a polymer might be modified to improve oxidative stability. The reaction of HO˙ radicals with a polymer electrolyte fuel cell membrane represents an initiation step for irreversible membrane oxidation. By means of pulse radiolysis, we measured k = (9.5 ± 0.6) × 10(9) M(-1) s(-1) for the reaction of HO˙ with poly(sodium styrene sulfonate), PSSS, with an average molecular weight of 1100 Da (PSSS-1100) in aqueous solution at room temperature. In the initial reaction of HO˙ with the oligomer (90 ± 10)% react by addition to form hydroxycyclohexadienyl radicals, while the remaining abstract a hydrogen to yield benzyl radicals. The hydroxycyclohexadienyl radicals react reversibly with dioxygen to form the corresponding peroxyl radicals; the second-order rate constant for the forward reaction is k(f) = (3.0 ± 0.5) × 10(7) M(-1) s(-1), and for the back reaction, we derive an upper limit for the rate constant k(r) of (4.5 ± 0.9) × 10(3) s(-1). These data place a lower bound on the equilibrium constant K of (7 ± 2) × 10(3) M(-1) at 295 K, which allows us to calculate a lower limit of the Gibbs energy for the reaction, (-21.7 ± 0.8) kJ mol(-1). At pH 1, the hydroxycyclohexadienyl radicals decay with an overall first-order rate constant k of (6 ± 1) × 10(3) s(-1) to yield benzyl radicals. The second-order rate constant for reaction of dioxygen with benzyl radicals of PSSS-1100 is k = (2-5) × 10(8) M(-1) s(-1). We discuss hydrogen abstraction from PSSS-1100 in terms of the bond dissociation energy, and relate these to relevant electrode potentials. We propose a reaction mechanism for the decay of hydroxycyclohexadienyl radicals and subsequent reaction steps.

Cysteamine reduces selenocystamine to form hemiselenocystamine and then cystamine. The rate constants are k(1) = 1.3 × 10(5) M(-1) s(-1); k(-1) = 2.6 × 10(7) M(-1) s(-1); k(2) = 11 M(-1) s(-1); and k(-2) = 1.4 × 10(3) M(-1) s(-1), respectively. Rate constants for reactions of cysteine/selenocystine are similar. Reaction rates of selenium as a nucleophile and as an electrophile are 2-3 and 4 orders of magnitude higher, respectively, than those of sulfur. Sulfides and selenides are comparable as leaving groups.