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PRODUCTION, DETECTION, AND CHARACTERIZATION

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[12] D i s t i n c t i o n b e t w e e n H y d r o x y l R a d i c a l a n d Ferryl Species

By J. D. RUSH, Z. MASKOS, and W. H. KOPPENOL Introduction

The Fenton reaction, or the reaction of ferrous ion with hydrogen peroxide at low pH, is generally considered to yield the hydroxyl radical. This radical is a strong oxidizing agent [E°( • OH/H20) = 2.73 V] 1 and attacks various small molecules with rates of 108-101° M -~ sec -1, whereas its reaction with proteins is diffusion controUed. 2 The products are carbon-centered radicals that in the presence of oxygen are converted to organic hydroperoxyl radicals) These radicals are rather oxidizing and can start various chain reactions. 4 The propagation reactions are well understood and are responsible for far more damage than the initiating event. It is clear that oxygen plays a dual role in this process: it makes the formation of the oxidizing species possible in the first place, and, second, it amplifies the damage through the chain reactions. While the concept of the hydroxyl radical as an initiator has received wide support, recent evidence suggests that at neutral pH and in the presence of a chelating agent a higher oxidation state of iron might be involved. This concept is not new: as early as 1932 it was proposed that a higher oxidation state of iron, the ferryl ion (FeO2+), might be involved in the decomposition of hydrogen peroxide. 5 While this appears not to be the case at low pH,6,7 the situation is more complex at neutral pH when iron is chelated. The failure of common hydroxyl radical scavengers to inhibit, for instance, the formation of ethylene from methionine resulted in the postulation of electron-donor or "crypto-. O H " complexes) Such a species should be fairly oxidizing to show more or less the same reactivity as the hydroxyl radical. For instance, a reduction potential of 1.2 V is ret W. H. Koppenol, Bioelectrochem. Bioenerg. 18, 3 (1987). 2 G. V. Buxton, C. L. Greenstock, W. P. Helman, and A. B. Ross, J. Phys. Chem. Ref. Data 17, 513 (1988). 3 K, U. Ingold, Acc. Chem. Res. 2, 1 (1969). 4 W. H. Koppenol, in "Oxygen and Oxy-Radicals in Chemistry and Biology" (M. A. J. Rogers and E. L. Powers, eds.), p. 617. Academic Press, New York, 1981. 5 W. C. Bray and M. H. Gorin, J. Am. Chem. Soc. 54, 2124 (1932). 6 C. Walling, Acc. Chem. Res. 8, 125 (1975). J. D. Rush and W. H. Koppenol, J. lnorg. Biochem. 29, 199 (1987); J. D. Rush, Z. Maskos, and W. H. Koppenol, FEBS Lett. 261, 121 (1990). 8 R. J. Youngman and E. F. Elstner, FEBS Lett. 129, 265 (1981).

METHODSIN ENZYMOLOGY,VOL. 186

Copyright© 1990by AcademicPress, Inc. All rightsof reproductionin any formreserved.

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DISTINCTION BETWEEN "OH AND FERRYL SPECIES

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quired to abstract an a-hydrogen from methanol, 9 and such a value would seem to be a lower limit. A thermodynamic derivation ~° suggests a value in excess of 0.9 V for the FeW/FeUI-EDTA couple. The term ferryl is commonly used to describe an oxidizing iron species derived from the reaction of hydrogen peroxide by ferrous complexes. Little is known directly about the structure and reactivity of high-valent iron in aqueous solutions, however, with the exception of that of ferryl porphyrins H,~2and some spectroscopic information and decay kinetics of ferryl and ferrate (FeV), FeO4 3-, in alkaline solution. 13Ferryl can be represented by Fe H (H202), F e II = O 2+, Fem 0% or Fe Iv ( O H - ) 2 , in which the formal charge is indicated by roman numerals. However, the oxidation state of iron is +4. Methods

The methods currently used for analysis of Fenton intermediates are summarized below. In large part they rely on discrimination between the anticipated effects of the hydroxyl radical and the experimental observations. However, even in systems of much less than biological complexity the analysis is not simple, and therefore experiments involving a minimal number of reactants are best. The formation of a reactive species is rate limiting and prevents its direct observation. Rate constants for the reaction of various ferrous complexes with hydrogen peroxide are given in Table I. 14-19

Inhibition of Peroxide Decomposition Catalytic decomposition of hydrogen peroxide takes place if a ferrous complex, HLFe II, is oxidized by peroxide and HLFe m is reduced by 9 W. H. Koppenol and J. D. Rush, J. Phys. Chem. 91, 4429 (1987). l0 W. H. Koppenol and J. Liebman, J. Phys. Chem. 88, 99 (1984). iID. Ostovic and T. C. Bruice, J. Am. Chem. Soc. 110, 6906 (1988). 12j. R. L. Smith, P. N. Balasubramanian, and T. C. Bruice, J. Am. Chem. Soc. U0, 7411 (1988). 13 j. O. Rush and B. H. J. Bielski, J. Am. Chem. Soc. 11)8, 523 (1986). 14j. D. Rush and W. H. Koppenol, in "Free Radicals, Metal Ions and Biopolymers" (P. Beaumont, D. Deeble, B. Parsons, and C. Rice-Evans, eds.), p. 33. Richelieu Press, London, 1989. i~ T. J. Hardwick, Can. J. Chem. 35, 428 (1957). 16O. K. Borggaard, O. Farver, and V. S. Andersen, Acta Chem Scand. 7,5, 3541 (1971). 17 H. C. Sutton and C. C. Winterbourn, Arch. Biochem. Biophys. 235, 106 (1984). is B. C. Gilbert and M. Jeff, in "Free Radicals: Chemistry, Pathology and Medicine" (C. Rice-Evans and T. Dormandy, eds.), p. 25. Richelieu Press, London, 1988. 19 S. Rahhal and H. W. Richter, J. Am. Chem. Soc. 110, 3126 (1988).

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PRODUCTION, DETECTION, AND CHARACTERIZATION

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TABLE I RATE CONSTANTS FOR REACTION OF FERROUS ION AND FERROUS AMINOPOLYCARBOXYLATE COMPLEXES WITH HYDROGEN PEROXIDE Rate constant (10 3 M -t sec -I) reported by Complex Aquo DTPA EDTA HEDTA NTA EDDA Fet~ATP FenADP Fe~citrate

HU

BFA b

SW c

RK d

GJ *

RRJ

41.5 - 0.3g ------

60 -+ 6g -9.10 -+ 0.08 16.7 -+ 0.1 18.4 +-- 0.1 --

-0.51 7 ----

-0.8 7 42 30 78 6.6 -+ 0.3 11 +-- 2 4.9 --- 0.3

200~ 13.5

1.37 _+ 0.07

10

° Hardwick ~5 (20.2 °, 0.1 N HCIO4); this reference contains a discussion of earlier work. b Borggaard et al. ~6(20 o, pH 0.2 M ionic strength); values apply to unprotonated species and were determined polarographically. ~ Sutton and W i n t e r b o u r # 7 (pH 7.4, varying ionic strength). d Rush and Koppenol 7 (25 °, p H 7.2 _+ 0.2, 38 m M ionic strength); stopped-flow study; Rush et al. 7 e Gilbert and J e f f 8 (pH 7); ESR study. f Rahhal and Richter ~9 (pH 7.0); rapid mixing study. Data in M t sec-i.

superoxide. This set of reactions [reactions (2)2(4)] is sometimes referred to as a metal-catalyzed H a b e r - W e i s s 2° reaction21: H L F e 111 + H202 ---, H L F e I~ + 02 ~ + 2 H ÷ H L F e II + H202 ~ H L F e O 2÷ + H20, or HLFe m + 'OH + OHH L F e O 2÷ + H202 ~ H L F e Ill + 02" + H20 • OH + H202 ~ 02 v + H20 + H + 02 ~ + H L F e m -o H L F e " + 02 H L F e O 2÷ + RH ~ H L F e m + O H - + R . • O H + RH--+ H20 + R"

(1) (2a) (2b) (3a) (3b) (4) (5a) (Sb)

2o F. Haber and J. Weiss, Proc. Roy. Soc. London A 147, 332 (1934). 2~ The reaction 02 = + H202 --~ • OH + O2 + O H - forms part of a cycle originally proposed by F. Haber and R. Wilst/itter, [Chem. Ber. 64, 2884 (1931)] to account for the decomposition of hydrogen peroxide, but it became known as the H a b e r - W e i s s reaction. The ability to react with hydrogen peroxide was considered an essential characteristic of the superoxide anion. It was s h o w n later by several investigators [see W. H. Koppenol, in "Oxidases and Related Redox S y s t e m s " (T. E. King, H. S. Mason, and M. Morrison, eds.), p. 127. Pergamon, Oxford, 1982] that the rate of this reaction was too slow to be an efficient source of hydroxyl radicals.

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151

Competition between reactions (3a) and (3b) and (5a) and (5b) reflects the relative reactivity of the hydroxyl radical and a putative hypervalent iron intermediate toward hydrogen peroxide and the organic scavenger, respectively. If the product of this reaction, R., does not propagate the chain reaction by reducing the ferric complex, a sequence of scavengers with varying reactivity toward the hydroxyl radical can be employed to determine if the relative rates of reactions (3a) and (5a) correspond to that of the hydroxyl radical. This system requires a substantial excess of hydrogen peroxide since its rate of reaction with the hydroxyl radical is low [k(. OH + Hz02) = 2 x 107 M -1 s e c -I] compared to most organic scavengers, including the commonly used ligands of iron(II)/iron(III). Quenching aliquots of reaction mixtures in a titanium sulfate-sulfuric acid solution is a convenient method to determine hydrogen peroxide concentrations to approximately 10-5 M accuracy if solution components do not interfere with the absorbance of peroxotitanium at 408 nm (e = 740 M -l cm-~), z2 Oxygen evolution may also be monitored with a Clark electrode as an index to the rate of peroxide decomposition. The oxidation of ferrous complexes by hydrogen peroxide generates 2 mol of iron(Ill) per mole of peroxide. The oxidizing intermediate may be reacted with excess of iron(II), giving the full stoichiometry, or with an added scavenger. The scavenger product, which is normally a radical, may react as follows: A [Fe(IIl)]/A [H202]limiting R ' + H L F e II + H + --~ H R + H L F e m R . + H L F e la ~ R + + H L F e 11 R" + R" ---) p r o d u c t s

2:1 0:1 1:1

(6) (7) (8)

In their reactions toward Fel~/FeUi-aminopolycarboxylate complexes, carbon-centered radicals are usually oxidizing when the site of hydrogen abstraction is a to a carbonyl or carboxylate group as in acetone or acetate, inert if adjacent only to alkyl residues as in tert-butanol, and reducing if a to the hydroxyl group of alcohols. If more than one kind of radical is formed, nonintegral limiting stochiometries are found. If R. reduces H L F e hI, the variation in stoichiometry of H L F e m production with the organic scavenger concentration should follow Eq. (9a), 7 where a.on = k(. OH + HLFen)/k( • OH + HR): AHLFeIII/AH202 = 2a.oH[HLFeU]/[HR] 1 + a.oH[FeUL]/[HR] 22 M. J. Irvine and I, R. Wilson, Aust. J. Chem. 32, 2283 (1979).

(9a)

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PRODUCTION, DETECTION, AND CHARACTERIZATION

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from which a.on can be compared with literature values. 2 For an inert radical R., Eq. (9a) becomes 7 a.on[HLFe n]/[HR] AHLFeIII/AH202 = 1 + I + a.oH[FeIIL]/[HR]

(9b)

Solutions must be thoroughly deaerated and the ferrous complex maintained in excess. The stoichiometry, and hence a, can be measured with good accuracy by the stopped-flow method, using the high absorptivity of iron(III) complexes in the UV region• A current limitation is that data on the rate of hydroxyl radical reactions with iron(II) complexes is limited. The method has been applied to iron complexes of EDTA, diethylenetriaminepentaacetic acid (DTPA), and (N-hydroxyethyl)diaminetriacetic acid (HEDTA). 7 We have used the argument that the inability of tertbutanol to scavenge the reactive intermediates formed from the interaction of a ferrous chelate with hydrogen peroxide is evidence for a higher oxidation state of iron. 23 The experimental observation is that 1.7 to 2.0 tool H L F e Ix is oxidized per mole of H202 consumed, and that tert-butanol does not decrease this ratio to 1.0 as expected for an inert tert-butanol radical [Eq. (9b)]. 7 An explanation has been offered by Gilbert and Jeff, ~8 who suggested that the tert-butanol radical is not inert, but oxidizes the ferrous chelate and undergoes a reductive elimination, as studied by Eberhardt,24 •CH~(CH3)2OH+ HLFOl + H+ ~ H2C~-C(CH3)2+ HLFem + H20

(10)

However, the amount of 2-methylpropene formed is small, and reaction (10) appears not to account for the observed ratio of 1.7-2 ferrous complexes oxidized per hydrogen peroxide consumed (J. D. Rush and W. H. Koppenol, unpublished, 1989). Rahhal and Richter ~9also studied the oxidation of ferrous DTPA by hydrogen peroxide in the presence and absence of tert-butanol and found less than expected scavenging, irrespective of the tert-butanol concentration. Product Analysis In Fenton-related systems, the oxidations of organic components by either the hydroxyl radical or a higher oxidation state of iron typically give radical intermediates. Direct determination of these organic radicals by EPR (electron paramagnetic resonance) flow methods or analysis of products can give misleading results if the effect of other solution components, particularly the iron complex, is not taken into account. For in23 j. D. Rush and W. H. Koppenol, J. Biol. Chem. 2,61, 6730 (1986). 24 M. K. Eberhardt, J. Org. Chem. 49, 3720 (1984).

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153

stance, the oxidation of alcohols in a FettEDTA-H202 flow EPR experiment detected only fl-carbon-centered radicals, 25 in contrast to the fact that hydroxyl radicals abstract hydrogen at the a position with about 90% efficiency. Since only the a-carbon radicals readily reduce FeIUEDTA, however, these results were not useful in assigning the intermediate to a selective iron oxidant. 6 A characteristic reaction of ferryl versus that of the hydroxyl radical in water has not yet been unambiguously identified. However, stopped-flow experiments showed that transient species absorbing around 300 and 410 nm are obtained when ferrous HEDTA, 7 nitrilotriacetate (NTA), and ethylenediaminediacetic acid (EDDA) react with hydrogen peroxide in the presence of formate ions. 26 Two plausible pathways for their formation have been suggested.

Pathway A26: H L F e n + H202 ~ H L F e O 2÷ + H20 H L F e O z÷ + HCO2- --~ • L F C I + CO27 + H20

(11)

Pathway B27: As in Eqs. (2b) or (2a) H L F e n + H202

~

• O H + O H - + H L F e m, or H L F e O 2÷ + H20

•OH + HCO~-

--* H20 + CO2~

H L F e O 2+ + HCO2- ~ H L F e m + CO2 ~ + O H COs ~ + H L F e II ~ HLFezzCO2 - or HLFemCO22-

(12a) (12b) (13)

In these systems, ligand degradation of EDDA and NTA by peroxide is not inhibited by sodium f o r m a t e , 26 which is consistent with one-electron oxidation of the ligand and not exclusively of formate ions.

Indicator Reactions The use of a stable scavenger product as a chromogen to monitor the generation of oxidizing transients is a valuable method that depends on the availability of species which give selective reactions and easily detectable and stable oxidation products. Usually the progress of hydroxyl or ferryl generation is monitored spectrophotometrically. As an example, the reaction of substoichiometric amounts of radiolytically generated hydroxyl radicals with ferricytochrome c leads via a surface reaction to reduction of the heme with approximately 50% yield. 28 Since noreduction 2~ T. Shiga, J. Phys. Chem. 69, 3805 (1965). 26 j. D. Rush and W. H. Koppenol, J. Am. Chem. Soc. 110, 4957 (1988). 27 S. Goldstein, G. Czapski, H. Cohen, and D. Meyerstein, J. Am. Chem. Soc. 110, 3903 (1988). 28 j. W. van Leeuwen, A. Raap, W. H. Koppenol, and H. Nanta, Biochim. Biophys. Acta 503, I (1978).

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PRODUCTION, DETECTION, AND CHARACTERIZATION

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was observed in the presence of ferrous EDTA and hydrogen peroxide, it was concluded that formation of hydroxyl radicals in this system is unlikely. 29 Three compounds of low molecular weight commonly used are p-nitrophenol (giving a red nitrocatechol), 3° 2,2'-azinobis(3-ethyl-l,2dihydrobenzthiazole 6-sulfonate) or ABTS 31 (giving the intensely green ABTS +" radical cation), and methyl viologen. ABTS *. is formed from the parent molecule by one-electron oxidation. The reaction is effected by hydroxyl radical (but with only 58% efficiency), 32 peroxidase, 3~ peroxytetrasulfenatophenylporphyrin iron(III), 33 and a variety of inorganic oxidizing radicals such as Br2 ~, which reacts with I00% efficiency. 32 Numerous other reagants, notably Fe 3+, can oxidize ABTS [E'(ABTS +./ABTS) ~ 0.6 V]; therefore, careful controls of separate components of a system should be performed prior to its use in Fenton-related systems. ABTS +. has useful absorptions at 415 nm (e = 3.6 × 104 M -1 sec -~) and in the region of 600 nm that can be used to monitor the generation of oxidants in systems which contain no added reductants. We have used this system to monitor oxidizing transients during the ferric EDDA-catalyzed decomposition of peroxide. Although ABTS can be added to a concentration that scavenges efficiently in Eq. (5a) or (5b), reducing transients such as ferrous complexes and superoxide often limit the net amount of ABTS +. formed. 26 Two observations by this method suggest that free hydroxyl radicals are not the exclusive oxidizing transients in the iron-NTA or iron-EDDA systems. 26 (1) The rate of ABTS-r. generation increases with ABTS at concentrations well above that requisite for complete hydroxyl scavenging. (2) The addition of bromide ions in any concentration does not promote ABTS .+ formation, indicating that bromide does not scavenge the species which oxidizes ABTS as expected for the following reactions: • OH + 2 B r - ~ Br2 ~ + OH Br2 ~ + A B T S ---~ A B T S '+ + 2 Br

(14) (15)

This suggests that a higher oxidation state of iron may be less reactive, for thermodynamic or kinetic reasons, than the hydroxyl radical toward bromide ions. The ABTS/ABTS .+ system is insensitive to oxygen but very sensitive to reducing agents. The radical cation is reduced by the hydrogen perox29 W. H. Koppenol, J. Free Radicals Biol• Med. 1, 281 (1985)• 30 T. M. Florence, J. lnorg. Biochem. 23, 131 (1985). 31 R. E. Childs a n d W. G. Bardsley, Biochem. J. 145, 93 (1975). 32 B. S. W o l f e n d e n and R. L. Willson, J. Chem. Soc., Perkin Trans• 2, 805 (1982). 33 M. F. Zipplies, W. A. Lee, and T. C. Bruice, J. Am. Chem. Soc. 108, 4433 (1986).

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DISTINCTION BETWEEN "OH AND FERRYL SPECIES |

tn

-1

w

!

155

i

F

-3

pH FZG. 1. First-order rate constants for reduction of the A B T S +. radical cation in 3.5 mM hydrogen peroxide (A) and as catalyzed by addition of 1.5 mM HLFe m, where HL is DTPA (B), NTA in phosphate buffer (C), NTA in MOPS buffer (D), EDTA (E), or HEDTA (F). The concentrations of the buffers were 10 mM (MOPS, Tris, or phosphate), and rates were dependent on the buffer only for NTA.

ide anion [ k ( H O 2 - + ABTS t) = 1.8 x 103 M -] sec-l], and this instability is enhanced by ferric complexes which bind with peroxide: HO2- + ABTS '* --~ ABTS + H ÷ + 02HLFe m + H202 ~- HLFem(O2H -) + H + HLFem(O2H -) + ABTS +. ~ HLFe nl + 02- + H + + ABTS 02- + ABTS -+--~ ABTS + 02

(16) (17) (18)

(19)

These reactions require that the iron complex used as a Fenton catalyst be fairly active so as to maintain a steady rate of ABTS .+ formation greater than its rate of decomposition. Some experimental data are given in Fig. l, in which the rate constants for reduction of ABTS +. are determined in aqueous peroxide solutions with different iron complexes. The rate of ABTS .+ reduction exhibited by each most likely is due to the strength of peroxide binding to the ferric complex. Solutions of approximately 2 × 10-4 M ABTS -+ can be prepared from ABTS by controlled potential electrolysis at an isolated anode. Continued oxidation leads to apparently polymeric products. These preparations should be used to check the stability of ABTS +"radicals with components

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PRODUCTION, DETECTION, AND CHARACTERIZATION

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of the experimental solutions before using the ABTS/ABTS +. system as an indicator.

Measurement of Free Radical Chain Length An approach employed by some workers involves the measurement of the chain lengths of free radical chain reactions. A known concentration of reductant is introduced into systems containing hydrogen peroxide, a metal (complex), and a scavenger such as formate 17,34 or methanol. 35,36 For instance, Sutton and Winterbourn 34 studied the following chain reactions, in which PQ indicates paraquat: pQ .+ + FenlEDTA ~ FelIEDTA + pQ2÷ FeUEDTA + H202 + H ÷ ~ FemEDTA + • OH + H20 • OH + HCO2- ~ H20 + CO2CO2 : + pQ2÷ __.>CO2 + PQ +'

(20) (2b)

(12a) (21)

Chain termination is thought to occur through the reactions: FeaEDTA + H~O2 ~ FeIIEDTA(H202) FeUEDTA(H202) + PQ' --->FemEDTA + PQ2+ + 2 OH-

(2a) (22)

and FelIIEDTA(H202) can also react with the formate anion. Reaction (2a) is thought to predominate, the ratio of hydroxyl radical versus higher oxidation state being 9: 1. Results have been presented to show that in the case of copper phenanthroline complexes that the propagation reaction is much slower than anticipated for a hydroxyl radical oxidant. 37 However, complicating factors such as cage reactions of the hydroxyl radical with the ligands might tend to obscure the interpretation. As in all the systems described, a complete knowledge of the reactions occurring among very reactive species is lacking. Acknowledgments This work was supported by a grant from The Council for Tobacco Research, I n c . - -

U.S.A.

34 H. C. Sutton and C. C. Winterbourn, Free Radical Biol. Med. 6, 54 (1989). 35 G. R. A. Johnson, N. B. Nazhat, and A. Saadalla-Nazhat, J. Chem. Soc. Chem. Commun., 407 (1985). 36 G. R. A. Johnson, N. B. Nazhat, and R. A. SaadaUa-Nazhat, J. Chem. Soc., Faraday Trans. 84, 501 (1988). 37 S. Goldstein and G. Czapski, J. Free Radicals Biol. Med. 1,373 (1985).

Distinction between hydroxyl radical and ferryl species.

148 PRODUCTION, DETECTION, AND CHARACTERIZATION [12] [12] D i s t i n c t i o n b e t w e e n H y d r o x y l R a d i c a l a n d Ferryl Species B...
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