Free-radical chemistry of sulfite.

The free-radical chemistry of sulfite oxidation is reviewed. Chemical transformations of organic and biological molecules induced by sulfite oxidation are summarized. The kinetics of the free-radical oxidations of sulfite are discussed, as are the kinetics of the reactions of the sulfite-derived radicals SO3 and the peroxy derivative SO5 with organic compounds.

and at very high concentration, disulfite.
HSO-+ S02eaq -t HS20O PKa = 1.5 (4) (3) At any physiological pH, sulfite and bisulfite will both be important forms of S(IV). We will use primarily the term sulfite to refer to the equilibrium mixture, except when referring specifically to bisulfite. The term S(IV) will be used to include other compounds containing sulfur in the + 4 oxidation state.
Sulfur dioxide can produce bronchoconstriction upon inhalation, particularly in asthmatics and during exercise (5,6). In addition to inhalation of SO2, sulfite can enter the body due to its use as a preservative in food, wine, and medications. Finally, sulfite is a likely intermediate in the metabolism of sulfur containing amino acids such as methionine and cysteine.
Both liver and lung tissues contain the enzyme sulfite oxidase which catalyzes the oxidation of sulfite to sulfate. This has led to two contrary views of the possible physiological consequences ofingested sulfite. One point ofview is that the body contains sufficient sulfite oxidase to detoxify any reasonably likely dose of sulfite from either inhaled atmospheric SO2 or from food additives (7). The other view is that sulfite reaches the blood and forms S-sulfocysteine, RSS03 and, therefore, the sub-*Chemical Kinetics Division, National Bureau of Standards, Gaithersburg, MD 20899. sequent chemistry of S(IV), at least as the S-sulfocysteine, must be considered (8). Further, epidemiological evidence suggests a relation between SO2 and lung cancer in workers exposed to arsenic and animal studies on benz(a)pyrene correlate cancer development with SO2 exposure (8).
Sulfite is a strong nucleophile and reacts with many biomolecules by substitution at electrophilic positions. These reactions have been reviewed by Petering (8) and will not be discussed here, other than the reaction of bisulfite with cystine [Eq. (4)].
RSSR + HSO-z± RSSO-+ RSH (4) This reaction has an equilibrium constant of 0.089 at pH 7.75 and 37°C (9). The large concentration of RSSR causes most sulfite in the blood to be bound as S-sulfocysteine, RSS03-. As Petering points out, the biochemistry of HS03becomes the biochemistry of RSS03 beyond the lung.
Because of the above equilibrium, however, S-sulfocysteine may act as a reservoir for sulfite; when it reaches cells in which RSH is in greater abundance than RSSR, e.g., liver cells, where RSH:RSSR = 102_103 (10), the equilibrium may shift to the left to produce sulfite.
The present review deals exclusively with elements of the radical chemistry of sulfite. In light of the discussion above, it might appear that radical reactions initiated by sulfite are likely to be unimportant. There are, however, two possible sources of radicals from sulfite that can be considered.
First, the lung and the rest ofthe respiratory system, being rich in oxygen, provide an environment for the autoxidation of sulfite before it can either be removed by sulfite oxidase or converted to S-sulfocysteine. The autoxidation of sulfite may be initiated by trace metal ions or certain enzymes and clearly involves free radicals (11,12). The second possible source of radicals is from S-sulfocysteine. The one-electron reduction of RSSO3 can be written as either RSSO-+ e-RS-+ SO § (5a) or RSSO-+ e--RS + SO2- (5b) Pulse radiolysis experiments in which cysteine radicals were produced in the presence of So32, or in which SO3radicals were R2roduced in the presence of cysteine, showed that RS oxidizes sulfite and that, therefore, the first path is more likely.

Chemical Transformations Induced by Sulfite Oxidation
Much of the interest in the chemistry of radicals derived from sulfite arises from the observations that the reaction of sulfite with several organic compounds requires the presence of an oxidizing agent, usually molecular oxygen. Complementary to these observations are the many studies that show that certain organic compounds inhibit the oxidation of sulfite by oxygen.
The investigation of the effects of organic substances on the rate of oxidation of sulfite solutions by oxygen was initiated by Bigelow (13) and carried on actively for several years (14,15). In the work involving the oxidation of sulfite catalyzed by trace metal ions, the inhibition could have been caused by the complexation of the metal ion. Therefore, studies were carried out in which sulfite oxidation was inititated by ultraviolet light (16). Again, organic substances were found to inhibit the reaction. The photochemical reaction was shown subsequently to be a chain reaction and the inhibition by organic compounds due to breaking the free-radical chains.
Since the inhibition of sulfite oxidation involves, in general, only small total amounts of reaction, products of the chain brealdng reaction usually have not been discussed. Also, in some cases the initial reactant might be regenerated in a secondary process. In other cases, however, the chemical transformation of the inhibitor was evident. This was observed initially for quinine sulfate and pyridine, which turned green, and hydroquinone, which became opalescent (16). Other work showed that the inhibition of sulfite oxidation by alcohols was accompanied by their oxidation (17). In subsequent work, the oxidation of sulfite in the presence of unsaturated compounds was found to result in the addition of sulfite to double bonds (18). With pyridine this leads to formation of N-pyridinium sulfonate (19). The reaction of hydroquinone with sulfite in the presence of oxygen is perhaps the most studied (20,21), since sulfite was used as a preservative in hydroquinone-based photographic developers (22). In this system two types of reaction appear to take place: (a) oxidation of the hydroquinone by sulfite radicals and by molecular oxygen, and (b) sulfonation of the quinone to form hydroquinone sulfonates (followed by oxidation ofthe latter to quinone sulfonates) (21).
From the point of view of this review, the most important observations have been on the transformation of biological molecules by sulfite in the presence of oxygen. Fridovich and Handler (23) have shown that a mixture of horseradish peroxidase, hydrogen peroxide, and a peroxidizible substance initiate sulfite oxidation. Indeed, they used the oxidation of sulfite as a sensitive test for the production of radicals in biological systems (24). Klebanoff (25) confirmed this finding and further reported that the oxidation of NADH by Mn2+, peroxidase, and 02 was stimulated by sulfite. Therefore, a biological system can initiate the oxidation of sulfite and the subsequent chain reaction can provide reactive intermediates capable of reacting with biological molecules.
Since this early work, there have been several papers on the oxygen induced reactions of biological molecules with sulfite. It has been found that oxygen is required for the complete sulfonation of protein S-H groups by sulfite (26). Sulfite was found to form sulfonates with 4thiouracil derivatives in the presence of oxygen and this reaction was observed to be inhibited by hydroquinone (27,28). Methionine has been shown to be oxidized to the sulfoxide in the presence of sulfite, 02, and Mn2+ (29). This reaction appears to be inhibited by superoxide dismutase. Sulfite cleaves DNA in the presence of 02 and Mn2+ (30), this reaction is inhibited by hydroquinone. The autoxidation of sulfite can destroy indole-3acetic acid (31) or tryptophan (32) and several nucleotides and nucleic acids have been shown to react with sulfite in the presence of oxygen (33).
Both n-carotene (36) and vitamin Bi (37) are destroyed during the autoxidation of sulfite. Finally, papain is inactivated during sulfite autoxidation in a reaction which leads to the incorporation of sulfite into the protein (38).
In this review, we will discuss the chemistry of the free radicals S03and S05-, key intermediates formed in the autoxidation of sulfite. In addition, we will discuss briefly the radicals S02 and S04and the ion HS05, due to their possible relationship to the behavior of sulfite in the body.

Formation and Detection of S03-Radicals
The sulfite radical is generally produced by the oneelectron oxidation of sulfite or bisulfite ions, either chemically or photolytically. The radical is detected either by ESR or by optical absorption spectroscopy. Although the ESR detection is more definitive, kinetic studies on the sulfite radical are more often carried out by absorption spectroscopy, by monitoring either the absorption of S03itself or more frequently by follo ing the formation of other more strongly absorbing sj cies arising from S03reactions with substrates.
The optical absorption of S03exhibits Xma,, = 2 nm with e_m = 1000 M-1 cm-' (48). This relativ4 weak UV absorption has been used to determine t second-order decay rate constant for this radical 2S03--S2026 or SO + SO32 -(2k=1.1x109 M-1sec-1) (3,44,49) but is not conve ently used for following the kinetics of S03 reactic with substrates since manx of these substrates or th radical products mask the SO3 UV absorption. The: fore, in pulse radiolysis experiments often the absoi tion of the other substrate radical was monitored.
,w-Kinetics of One-Electron Oxidation peof Sulfite of As mentioned above, sulfite or bisulfite ions undergo ion 9ne-electron oxidation by several radicals to produce re-S03-. Rate constants for a number of reactions of this of type have been determined by pulse radiolysis and are by summarized in Table 1. The hydroxyl radical reacts with or both sulfite and bisulfite with very high rate constants, near the diffusion-controlled limit. The rate of oxidation by other radicals decreases in an order that appears to reflect the order of expected oxidation potentials of (6) these radicals. Measurements of rate constants over a wide range of pH allows the separate determination of igh rate constants for the oxidation of sulfite and bisulfite.
Whereas the hydroxyl and sulfate radicals react with bisulfite about twice as fast as with sulfite, for every other radical the reaction with sulfite is the faster by (7) far. For Br2 the ratio is about 4; for the weaker oxidant I2 , the ratio is about 200. For the dimethylaniline rad-1).
C6H5f4H+ + HS03--C6H5NH2 + H+ + S03-(12) C6H5I4H + SO32(HSO-) --no reaction (13) lue The neutral aniline radical, C6H5NH, on the other hand, does not oxidize sulfite. The cation radicals from promethazine, tryptophan, and tryptamine also oxidize HS03 with moderate rate constants (Table 1) and in (9) these cases the reactions were found to lead to equilibrium. The reverse reactions and equilibrium constants ar-will be discussed below. ex-Since the autoxidation of sulfite solutions was found nd to be catalyzed strongly by trace amounts of transition al-metal ions, the reactions of sulfite with metal ions in the their higher oxidation states has been the subject of as many studies (11). Frequently, these studies are complicated by the ability of sulfite to complex these metal )55 ions. These complexes are often quite stable; mercuric ely ion ( in the presence of chloride) is used to protect sulfite the from air oxidation (67). Other metal ion-sulfite complexes are more labile, decomposing presumably to the reduced metal ion and the sulfite free radical. For strong oxidants like Mn(III), the reaction is fast and apparently (10) irreversible (68). For weaker oxidants like Fe(III), the reaction is much slower and reversible, making the deni-rivation of an elementary rate constant for the oxidation rns of sulfite difficult. Leir For substitution-inert metal ion complexes, the sitre-uation is somewhat simpler since complex formation by rp-sulfite is not important. Rates have been measured for the reactions of several metal ion complexes and rate HP04-+ So32 9 2.7 x 107 (60) po42-+ SO32- I2-+ HSO37/SO32-6.7 1 x 107 (63) I2-+ So32- 11 1.9 x 108 (63) NH2 + So32- -5 x 108 (59) (promethazine)Y+ + HS03-3.6 6 x 108 (59) (promethazine)Y + HSO3-iSO32-6.6 1.2 x 108 (59) (tryptophan)Y+ + HS03-  (66) aNo reaction detected (k<108M-lsec-1). The redox potentials for NH2 and S03 radicals appear to be very similar, judging from rate constants for their reactions with several reactants. b Calculated from the pH dependence of the rate constant.
constants derived for the primary step, the one electron oxidation of sulfite. Bisulfite and S02 aq are assumed to be unreactive and the reported values depend upon the acid dissociation constants chosen. Some values reported are given in Table 2.
There is no information on the oxidation by free radicals of S02 aq, the form of sulfite present in strong acid. Mn(III), a strong oxidant (E=1.4 V), does react with S(IV) in 2-6 M HC104. The very strong inverse dependence of the rate constant on the acid concentration was interpreted as showing that bisulfite was the important reactant, not SO2 * aq (68).

One-Electron Reduction of S(IV)
Although not of importance in autoxidation, the reduction of S(IV) could be important in some biological systems. The hydrated electron is unreactive toward So32-(k < 106 M-1sec-I) and reacts very slowly with HS03to produce hydrogen atoms (k = 2 x 107WM-1sec-1) (3). On the other hand, SO2 is reported  (74) to be reduced rapidly by C02 to produce S02-, while HS03and So3 were unreactive (75). S02also is produced by the reduction of bisulfite using methyl viologen radical, flavodoxins, and in a H2/hydrogenase system (76, Z7). More recently, enzymatic reduction of bisulfite to S02 was demonstrated in hepatic microsomal protein and ascribed to reaction of cytochrome P-450 (78). Also Ti3+ was found to react with sulfite in acid solutions (pH 2-6) to produce S02 (42). All the above reactions probably occur by electron transfer to S02 rather then bisulfite. The radical S02-produced in these processes is in equilibrium with dithionite (S2042-) and is known to be a highly reactive one-electron reductant. It reduces metalloporphyrins containing Fe(III), Co(III), and Mn(III) (79-81) and a wide variety of electron-transfer proteins (82). The reactivity of S02appears to follow the same pattern as 02, with rate constants about 103 times higher (83). The potential for the process S02(aq) + e -8 02 (14) has been estimated as -0.26 V (84).

Reactions of Sulfite Radicals
The S03radical is for the most part a sulfur-centered radical which can act as an oxidant or reductant and like most other radicals may engage in hydrogen abstraction or addition to double bonds. Hydrogen abstraction, e.g., from isopropanol, was found to be unimportant (k 103 M1 sec) (3). This finding is not surprising, since the S-H bond expected to be formed in this process is much weaker than the C-H bond. (Formation of an 0-H bond on sulfite is not likely due to the low spin density on the oxygens of this radical) (56).
Addition of sulfite radicals to unsaturated bonds (C = C, C = N, and C-C) has been demonstrated by ESR (39)(40)(41)(42)49,55). These reactions were found to be very sensitive to steric effects by substituents on the unsaturated bond. Because of the steady-state nature of these ESR experiments no kinetic data are available. Attempts to measure addition rate constants by pulse radiolysis using allyl alcohol as an example gave only an upper limit of 106 M-1 sec-' (66). The ESR results demonstrate, however, the feasibility of sulfite radical addition to unsaturated biological targets. Extensive kinetic studies were carried out by pulse radiolysis on the oxidation of organic substrates by 503 . The results are summarized in Table 3. The sulfite radical is found to oxidize ascorbate, trolox ( a watersoluble tocopherol derivative), methoxyphenol, hydroquinone, phenylenediamines, and chlorpromazine with moderate rate constants varying in the range of 106 to 109 M-1sec-1, depending on the redox potential of the substrate and on the pH, for example with ascorbate S §0 + H2A -HSO + A-+ H+ k<106 M-lsec'- (15) S03 + HA-SO2-+ A-+ H+ k=9x106M-1sec-1 (16) SO-+ A2-S02-+ A-k=3x108 M-'sec- (17) For hydroquinone, catechol, and several other di-and trihydroxybenzenes, the effect of pH on their reactivity with S03 was demonstrated in detail (64). All of these compounds were unreactive in neutral solutions but became highly reactive as they deprotonated in basic solutions. Compared to other oxidizing radicals such as Br2 , I2- (86), and phenoxyl (87), S03 reacts more slowly and appears to be a milder oxidant. From a redox equilibrium established between bisulfite and chlorpromazine at pH 3.6 [Eq. (18)], S03 + ClPz + H+ a HSO + ClPz + (18) the redox potential for the couple S03-/HS03was measured to be 0.84V vs. NHE (59). The redox potential for the S032 /SO32 couple in basic solutions is calculated (from the pKa of HS03-3 SO + H+) to be 0.63 V vs. NHE. This change in potential explains why So32is oxidized by the same oxidant more readily than HS03 , as discussed above (Table 1).
Since the SO3-/SO3 -potential is now known, reactions of S03can be used to determine the potential for the one-electron oxidation of other species, in those cases where the electron transfer reaction is fast enough so that the decay of 503 due to self-reaction is not important. This was initially carried out for phenol (59), leading to a new value of its one-electron redox potential. More recently, equilibrium constants also have been measured for the reactions of 03with tryptophan, tryptamine, and tryptophanamide (65), and dimethylaniline (63).
Knowing the redox potential for the reduction of S03allows us to calculate its oxidation potential from the known two-electron redox potential for So32in basic solution: S04 + H20 + 2e--S032 + 20H-E = -0.92 V (19)  Subtracting E(SO3-) = 0.63 V from twice the former value (-0.92 V) leads to S042 + H20 + e e~S03 + 20H E = -2.47 V (21) This suggests that 03can act as both a mild oxidant or a strong reductant. It may be difficult to demonstrate the reducing power of S03since many oxidants will react with sulfite ions before the 503radicals are produced in the radiolysis. Moreover, the above calculation of redox potential may not reflect the actual reducing power of S03since the initialTproduct is SO3, which is subsequently hydrated to S04 , possibly much more slowly than the electron transfer (as argued for the case of S02-) (84). In addition, it has been argued on the basis of spin density on the sulfur, that SO3 is a much weaker reductant than 0O2 (55).
Biological damage by S03 may be partly due to oxidation reactions similar to those in Table 3. But the main harmful effects of this radical may lie in the fact that it reacts very rapidly with 02, k = 1.5 x 109 M-1sec-1, to form a peroxyl radical which is much more reactive. * + 0-s°3 -+ 2 >°( 22) The alternative reaction path forming SO3 + O2was 213 .0-+ e---* S03'-found (85) to be unimportant, at least under the experimental conditions of pH 3-12.
The one-electron reduction of S05-yields HS05-, peroxymonosulfate (Caro's acid). This is a strong oxidant, with a standard two-electron reduction potential of 1.82 V (88).
HSO-+ 2H+ + 2e--HS04 + H20 (27) Peroxymonosulfate is known to oxidize many organic compounds (89). Of considerable interest is its ability to oxidize sulfides to sulfones (90) and primary aryl amines to nitroso compounds (91). In addition to these reactions with organic compounds, which might involve oxygen atom transfer, peroxymonosulfate can be reduced by metal ions, ppssibly producing the highly reactive free radicals, S04 or OH, as it does upon reaction with e,q- (92), e.g.
The S65radical possibly can also react by atom transfer. This mechanism has been proposed for its reaction with bisulfite (30).
If SO5is capable of transferring an oyxgen atom, the direct oxygenation of qrganic compounds is feasible, possibly producing the S04radical as a by product.
The existence of this type of reaction has not been confirmed.
The mechanism of the self-reaction of SO65is not known. It has been proposed that the reaction leads to stable products and serves to terminate the autoxidation of sulfite (11). On the other hand the reaction has been proposed to go to S04or S208 (98).
2SO6 -* 2SO4 + 02 2SO6 -* S2o°+ 02 (31) (32) with the ratio k(SO2-)/k(S20 2-) = 9. Although this reaction is not likely to be important in the physiological role of sulfite, it could be important in some of the laboratory studies of the effects of sulfite oxidation on biological systems. The mechanism of this reaction certainly deserves more study.

Reactions of Sulfate Radical
The S04radical, possibly produced by the reactions discussed above, is very reactive toward organic compounds. It can abstract H atoms, add to double bonds, and oxidize by electron transfer quite rapidly. The rate constants for such reactions with many organic compounds are summarized in a recent compilation (86) (see examples in Table 4) and will not be discussed here. It is clear, however, that S04 attacks biological targets indiscriminately.

Conclusions
It has been apparent for some time that the effects of SO2 autoxidation on organic and biological systems are due to reactive internediates. Since the reactivities of many of these intermediates are now known, the mechanism of these effects can be better understood. For several of the organic compounds, like hydroquinone and other phenolic species, reaction with S03and S65is possible. Indeed, they prove to be the most efficient inhibitors of SO2 autoxidation. For other organic compounds, like mannitol and methionine, only reactions with S04 are, likely. Production of HS05 from the reduction of S05 opens up additional possibilities, including direct reaction of HS05 and its decomposition to produce S04 or OH.
Within the body, it is apparent that if SO2 is allowed to undergo autoxidation, cellular damage is inevitable. Whether the presence of S(IV) beyond the region of the lungs can lead to similar damage is not apparent. To a large extent this damage will depend on the equilibrium RSSR + HSO-a± RSSO-+ RSH and the probability of forming radicals from RSS03-. One-electron reduction of this compound is expected to yield SO3radicals. This was recently supported by pulse radiolysis experiments, whereby reduction of RSS03 by eaqand (CH3)2COH was found to form a radical which oxidized ascorbate with the same rate constant as does S03- (66). If reduction of RSS03to SO3radical occurs with biological reductants, this could lead to oxidative damage by the S03and the other radicals produced from it. Thus the RS group serves not only as a carrier of sulfite but it also changes the requirements for S03radical formation from oxidation to reduction.