Summary
Free radicals are important intermediates in natural processes involved in cytotoxicity, control of vascular tone, and neurotransmission. The chemical kinetics of free-radical reactions control the importance of competing reaction pathways. Equilibria involving protons often influence the reaction kinetics of free radicals important in biology. Radiolysis is a powerful method to generate specific free radicals and measure their reactivity. Current work in this area at the CRC Gray Laboratory is providing information fundamental to our understanding of the molecular biology of cytotoxic and physiological processes, and leading to the identification of new targets for exploitation of cellular chemistry.

Introduction
A common misconception amongst the general public is that `natural' foods are always `good for you'. Extending this logic to biological processes, this raises the question as to how the body can destroy or otherwise eliminate unwanted molecules or organisms, using the normally apparently harmless molecules at its disposal. In a review in an earlier Annual Report I described how chemicals called free radicals derived from drugs can be designed to act as `magic bullets' in cancer chemotherapy and diagnosis, to help kill or mark specifically cells lacking oxygen in tumours (Free radicals: magic bullets in cancer therapy and diagnosis? Gray Lab. Ann. Rept., 1990, p.13; copies may be obtained from the author). This article concentrates on the chemical reactions, involving free radicals, the body makes use of to respond to diverse challenges, reactions which may result in unwanted injury if the natural defences are overwhelmed.

This is an area in which research has grown enormously in recent years, free-radical research now possibly rivalling radiation research in terms of resources devoted world-wide. Chemical bonds are usually formed from the sharing of two electrons, whereas a free radical is a species with one unpaired electron. This makes many, but not all, free radicals chemically quite reactive, as the species seek to find another electron to pair up with. However, the definition includes common chemicals such as oxygen. Not surprisingly, therefore, oxygen is a common reactant in free-radical processes, having a propensity to take part in single-electron transfer or free-radical addition reactions in which electrons become paired. Another common gaseous chemical which is a free radical is nitric oxide. It is now recognized to play a critical role in vascular physiology, and with its molecular formula of NO, this has led to as many puns in reviews of its role as the diverse roles themselves. Apart from the title of this article, the mind soon turns to phrases such as `NO sex, please, ...' (since nitric oxide is involved in penile erection). Biological messengers are often needed to be short-lived, degradable, controllable and reusable: the superoxide radical, the electron-adduct of superoxide, fulfils all these criteria.

This article outlines some of the chemical background to the use nature makes of free radicals in ordinary biological processes. The challenge in cancer research is to find ways of exploiting this chemistry for therapeutic gain, and some of our exploratory approaches are mentioned briefly in the report of the Molecular Mechanisms of Therapy Group. In contrast to the reductive free-radical processes stressed in the earlier Magic bullets ... review, free-radical reactions the body uses naturally are often oxidative in nature, as described below. To help keep the reactions in check, antioxidants to mop up unwanted free radicals have evolved, such as vitamin C (ascorbate) or vitamin E. Glutathione, a natural thiol commonly involved in coupling reactions to help eliminate unwanted chemicals by renal excretion, for example, also plays a part in the action of these antioxidants.

Ionizing radiation ejects single electrons from molecules, and so the radiolysis of water, for example, generates free radicals at a rate readily controlled by manipulating the radiation source. By adding suitable solutes, specific free radicals identical to those the body produces naturally can be generated. Thus techniques originally developed to help study the reactive intermediates in chemical events following radiolysis have proven to be powerful methods of characterizing the reactions of natural free radicals in predominantly aqueous media such as the bulk of the cellular environment. In fact, the rates of several of the most important reactions of natural free radicals were first measured in this way, and the tools of radiation chemistry provide important information concerning free radical reactions of biological importance. Radiation chemistry is therefore becoming increasingly important in this much wider context.

Free radicals and cellular oxidative stress
Oxidation and reduction are chemical terms which describe the loss or gain of electrons by molecules, respectively. Thus ferrous iron (Fe²⁺) is oxidized to ferric iron (Fe³⁺) by the loss of a single electron, the charge on the ion changing from +2 to +3 in the process. Hydroxide ions in water (OH^- ) can be ionized, losing an electron, to give hydroxyl free radicals (·OH); the unpaired electron in ·OH is denoted by the radical `dot', and such species have a strong tendency to restore the electron pair by pulling a hydrogen atom, complete with a single unpaired electron, from C-H bonds in sugars, in DNA for example. An oxidizing agent is thus a molecule, atom or radical fragment which likes to gain an electron. (In fact, radiation chemistry has provided us with a versatile and powerful method of quantifying the propensity to electron gain or loss involving short-lived free radicals, where conventional electrochemical methods fail.)

In the present context, oxygen itself is one of the commonest oxidizing agents. When oxygen acts as an oxidizing agent, it gains one or more electrons from a substance. If it adds a single electron, the superoxide free radical is formed (O₂·^- ). This is an extremely common substance being produced in our bodies all the time: it has been estimated that up to 2% of the oxygen used in mitochondrial respiration could end up as superoxide, although the figure may be less in healthy tissue. Thus oxygen is a common terminal electron acceptor in biochemical processes. Superoxide radicals are also a key feature of the phagocytic process (see below). Oxygen itself is a free radical, but one with two unpaired electrons; reduction by adding one electron to give superoxide involves the pairing of two of the electron spins, leaving one unpaired. (Electron spin is a property seemingly discovered by physicists to make chemistry more complicated, although the property is put to good use in detecting free radicals, just as nuclear spin is now widely used as the basis for magnetic resonance imaging.)

The involvement of superoxide radicals in biochemical processes has been accepted for only about 20 years, but some time before that, radiation chemists had recognized it was produced in irradiated, aerated (oxygenated) water. The key properties of superoxide radicals had thus already been mapped out and the discovery of an enzyme specifically reactive to superoxide (superoxide dismutase, SOD) further stimulated radiolysis studies of the chemical properties of O₂·^-. It was apparent that superoxide was not itself very reactive, at least compared to the rates of many free-radical reactions, which often occur immediately on collision of the reactive species (`diffusion- controlled'). Superoxide might thus be viewed as like a pro-drug, not itself directly damaging but leading to the production of a species much more reactive. The emergence of nitric oxide on the physiological scene has added a new twist to the pivotal role superoxide now seems to occupy in diverse biological processes. The chemical background to the consequences of superoxide production in biology, in particular how superoxide is converted to more damaging species, is the main theme of this article.

Superoxide on its own+
Superoxide is the dissociated form of a weak acid, the hydroperoxyl radical, HO₂·. Just like other weak acids, such as acetic acid (CH₃CO₂H , vinegar), HO₂· dissociates in water:

Thus when superoxide is formed in water by adding an electron to oxygen, there is a very rapid equilibration according to equation (2) and the concentration of protons (H⁺), i.e. the pH or acidity of the solution. The proportion of HO₂· to O₂·^- depends on the pH and the equilibrium constant of reaction (2), K₂.

(The equilibrium constant, K of a reaction is effectively the product of the concentrations of the products of the reaction, on the right hand side of the equation, divided by the product of the concentrations of the reactants, on the left. Thus for reaction (2), K₂ is defined as [O₂·^- ][H⁺][HO₂· ], where square brackets denote concentrations. The ratio of concentrations of O₂·^- to HO₂· is given by K₂ / [H⁺]: the more acidic the solution, the lower the fraction of the radicals in the O₂·^-   form. The higher the value of K the more the equilibrium will procede to the right as written; a value much less than unity implies the reaction is unfavourable. For mathematical convenience the equilibrium constant K is often described in the same way as the proton concentration is described, i.e. as pK₂. Comparison may be made with pH; `p' in both shorthand notations is a mathematical operator, equivalent to writing `-log₁₀'. Thus pH is essentially the same as -log₁₀[H⁺] and pK the same as -log₁₀K.)

The concentration of HO₂· equals that of its `partner', O₂·^-   when the pH has the same value as pK₂. Coincidentally, K₂ is close to that of pK₁ (hydroperoxyl is about as weak an acid as acetic acid), pK₂ being about 4.8. This value means that only at around pH 5 is the concentration of superoxide radicals in the HO₂· form about equal to that of O₂·^-. At physiological pH values close to 7, only around 1% of superoxide radicals are in the hydroperoxyl (HO₂· ) form.

This distinction between HO₂· and O₂·^- is important for two reasons. Firstly, the negative charge on O₂·^- will inhibit its diffusion across lipid membranes, since charged species generally have much lower solubility in lipids than uncharged molecules. Secondly, equilibrium (2) controls the `natural' lifetime of superoxide radicals in the absence of any other reactant or catalyzing enzyme. Radicals often react with each other, a natural tendency to pair up their electron spins. Superoxide is no exception, and the overall reaction between two superoxide radicals can be written as:

Since superoxide in water will always be subject to equilibrium (2), equation (3) is, in fact, properly described by three separate reactions yielding common products:

Radiation chemists measured the rate of reaction of superoxide/hydroperoxyl radicals as they reacted with themselves. The decay was pH-dependent, and was decreased as the pH increased. At very high pH they found no evidence that reaction (3c) occurred at all, so far as could be measured, and so superoxide is a stable species in the absence of protons (at very high pH). A solution of superoxide radicals in strong alkali can be kept in the refrigerator overnight if metal ion impurities are absent. Because many reactions of free radicals are often very fast, their study has sometimes been labelled `fast reaction chemistry'. Clearly, this is a misleading label.

Radiation chemists showed that the `natural' lifetime of superoxide radicals increased by a factor of ten per pH unit with increasing pH above pH 6, as could be predicted from the mathematics if reaction (3c) does not occur at a measurable rate, pK₂ « 6, and the decay of O₂·^-   must always involve a proton. (It is noteworthy that the lifetimes of free radicals produced on one-electron reduction of nitroimidazole or benzotriazine-N-oxide bioreductive drugs follow exactly the same type of decay kinetics as superoxide, i.e. the radicals only react with themselves in the presence of protons and so their lifetimes, like superoxide radicals, are pH-dependent.)

Superoxide radicals will thus tend to disproportionate or `dismutate' to hydrogen peroxide and oxygen. Measurements of reaction rates - chemical kinetics - of possible competing reactions enables us to predict the fate of superoxide radicals, i.e. whether their production leads to an increase in hydrogen peroxide or to some other species. For example, in competition to uncatalyzed disproportionation of superoxide (reaction (3)), superoxide may react with metal centres in enzymes such as superoxide dismutase, or with free iron. Reaction with the enzyme accelerates the removal of superoxide and the formation of hydrogen peroxide, and reduces the chance of reaction with free iron. Both these reactions are discussed further below.

While hydrogen peroxide is itself an oxidizing agent, it is the combination of hydrogen peroxide and superoxide which yields a much more reactive oxidizing agent the hydroxyl radical.

Haber-Weiss sixty years on
Fritz Haber (the discoverer of the nitrogen fixation reaction to make ammonia for fertilizers from nitrogen and hydrogen) and Joseph Weiss (who was later to become a distinguished radiation chemist working in Newcastle) proposed around 1933 that hydroxyl free radicals (·OH) were produced when superoxide and hydrogen peroxide react together:

The hydroxyl free radical is important in radiobiological damage and is several orders of magnitude more reactive towards cellular constituents than superoxide radicals (and many orders more reactive than hydrogen peroxide). Reaction (4) thus attracted a great deal of attention as a potentially-important way in which the cell could generate, in the absence of ionizing radiation, highly-reactive and damaging hydroxyl radicals. However, the quantitative methods of radiation chemistry continued to play a definitive role in the development of ideas about the mechanisms involved in cellular oxidative stress. Radiation chemists compared the rates of reactions (3) and (4), and concluded that reaction (4) was too slow to be important in biology, especially considering the acceleration of (3) by an enzyme, superoxide dismutase, then recently discovered (see below). This underlines the importance of considering not only the products or course of a reaction, but also its rate; pulse radiolysis and other radiation-chemical methods are by far the most powerful methods of obtaining rate information about chemical reactions involving free radicals.

Attention soon switched to other ways in which superoxide could lead to more damaging species. About 100 years ago, Henry Fenton (working in Cambridge) had observed that the reducing agent, ferrous iron (Fe²⁺), together with hydrogen peroxide could oxidize some organic compounds. Partly as a result of studies in the 1940s by the distinguished British radiation chemist, John Baxendale, the mechanism is now known to involve hydroxyl radicals, with a key step analogous to reaction (4) but with the electron donor, O₂·^-   replaced by Fe²⁺:

Since superoxide can act as an electron donor to suitably electron-affinic molecules, it was recognized that it could act to reduce the ferric iron produced in reaction (5) back to ferrous, thereby `cycling' it. This enables only trace levels of iron to catalyse the formation of potentially large quantities of hydroxyl radicals from superoxide. Reactions (5) and (6) effectively sum together to yield the same result as (4) but it is achieved much faster:

The basic chemical reactions by which univalent reduction of oxygen in the cell could lead to the generation of powerful and damaging oxidizing species were now established. What is still much less clear is how these reactions are controlled by the cell, or even if these are the main reactions underlying the need for the cell to evolve methods of generation or modulation of the concentration of superoxide radicals. Although reactions (4) to (6) have been quoted in countless studies in the last decade or so, their importance may turn out to have been overestimated: chemical kinetics can provide the key information to evaluate these questions, and radiation chemistry has provided both tools and experience directly relevant.

Iron and copper: controlling the heavy mob
The copper redox states (cupric/cuprous, Cu²⁺ /Cu⁺ ) can function in much the same way as iron in catalyzing the Haber-Weiss reaction (4). Thus the intracellular levels of free heavy metals, particularly copper and iron, are critical in defining the extent of hydroxyl radical production from superoxide and hydrogen peroxide. Except in some disease states, the iron storage protein, ferritin, maintains the iron tightly bound so that the free levels are very low. Of course, it is not just the average concentration of free iron (or copper) that is important, but where within the cell it is found. Thus the chemotherapeutic agent, bleomycin, is thought to exert DNA damage by binding iron. (Very powerful iron binders, or `chelators', may either complex the minute amount of free iron in the cell, or effectively `pull it off' weaker iron/protein complexes.) Since the bleomycin/iron complex binds to DNA because of other features of its molecular structure, any hydrogen peroxide in the immediate vicinity of DNA is immediately converted to hydroxyl radicals, which can cause a localized high concentration of strand breaks. Thus:

To generate hydroxyl radicals from hydrogen peroxide (reaction (5)), iron has to be in the reduced state. Although superoxide can reduce ferric iron, including liberating free ferous iron from iron bound in ferritin, it is transported in the ferric form by the protein, transferrin, which is resistant to reduction. Both ferritin and transferrin bind iron tightly, so that the levels of free iron are very small. They can be reduced even further by artificial chelators, which have been shown to reduce the effects of reperfusion in ischaemic heart disease (see below) in animal models. This raises the question of whether iron- supplemented diets are wise except in obvious anaemia; it has even been speculated that iron-catalyzed oxidative stress via the free-radical mechanisms discussed above is the basis for the sex difference in predisposition to heart disease, since serum ferritin levels are much higher in men than women.

Although copper, if free in a reducible form, can enhance free-radical damage, its most important role in controlling free-radical reactions is as a redox centre in the enzyme, superoxide dismutase (SOD). This catalyses (speeds up) reaction (3), in a manner which can be represented by:

Here Cu(II) and Cu(I) represent cupric and cuprous forms of the copper redox site, respectively. Superoxide acts first to reduce Cu(II), and then to oxidize Cu(I) (the species O₂^2- in reaction (9) is, in fact, effectively hydrogen peroxide since it immediately reacts with water/protons just as hydroxide is largely neutralized to water at pH 7). The activity of the enzyme and the rate of turnover was beautifully demonstrated in some pulse radiolysis experiments by Martin Fielden and colleagues soon after the discovery of the enzyme, about 20 years ago; it is extremely easy to produce superoxide radicals in a microsecond or so by pulse radiolysis.

Making more superoxide by ischaemia followed by reperfusion
Little mention has been made so far of the biochemical sources of superoxide. Numerous studies have been made of the enzyme, xanthine oxidase. This usually exists mainly in a form which catalyses the oxidation of xanthine and hypoxanthine to uric acid, using nicotinamide adenine dinucleotide as the electron acceptor. However, in tissues where the oxygen supply becomes used up following ischaemia (as in myocardial infarction, stroke, and in some tumour cells) the enzyme is changed to a form which has the potential to reduce oxygen to superoxide. On re-admission of oxygen (e.g. by reperfusion), there is a burst of radical generation. This is now thought to cause much of the damage in ischaemia-related diseases and in tissues for transplant. Species differences have made studies of these processes difficult. Thus it it not the shortage of oxygen itself that is damaging, but the reintroduction of oxygen after the cell has undergone biochemical changes caused by its absence. Again, the site of free-radical generation is likely to be important, and so there is much interest in identifying the types of cell (e.g. endothelial cells) where xanthine oxidase activity might be highest - or most important because of site-localized iron, for example.

Redox cycling: the flip side
In the earlier review on bioreductive drugs, the basis for the selective toxicity towards hypoxic cells was ascribed to activation of the drug to a free-radical intermediate reactive towards oxygen. The drug was thus only effectively reduced to a toxic product at low oxygen concentrations, i.e. in hypoxic cells. In oxic cells, the futile redox cycling (alternate acceptance and donation of an electron by the drug) served as a protective mechanism:

Obviously, reaction (11) amounts to a stimulation of cellular superoxide production, another potential toxin. Thus, whether it is truly protective depends on which is the lesser of two evils: more superoxide or the product(s) resulting from drug reduction.

The dirty bits
Sorry, I'm not really going to discuss the chemical basis for penile erection: this is a family publication. Let's just say `brewer's droop' can be induced by blocking an enzyme producing the free radical, nitric oxide. In fact, the control of vascular tone by free-radical `messengers' has been recognised only recently. Around 1980 it was demonstrated that the vascular endothelium (the cells lining blood vessels) controlled vascular tone by the release of a then unidentified substance labelled `EDRF': endothelium-derived relaxing factor. About seven years ago, EDRF was identified, to everyone's astonishment, as the simple free radical, nitric oxide (NO ). A critical experiment was the inactivation of EDRF by superoxide, and the enhancement of its biological activity by superoxide dismutase. This was ascribed to the radical-radical reaction:

in which peroxynitrite (ONOO^- ) was formed. This reaction was known previously, but its rate could not be established until pulse radiolysis and flash photolysis provided direct measurements.

Peroxynitrite, like superoxide, is stable in strongly alkaline solution. However, at physiological pH its lifetime is of the order of a second. This is because peroxynitrite is the salt or ionized form of the weak acid, peroxynitrous acid (HOONO), and the pK_a for its dissociation (ionization) is about 7:

Thus if peroxynitrite is formed at physiological pH, almost half equilibrates extremely rapidly to peroxynitrous acid. The latter rapidly rearranges to nitrate (NO₃^- ) in water, in competition with a reaction which has been thought to involve decomposition to hydroxyl radicals and nitrogen dioxide (NO₂):

Whether free ·OH is really formed, as shown in (15), is the subject of much current debate; certainly, peroxynitrous acid is a powerful oxidant with cytotoxic properties. It can initiate lipid peroxidation and thiol oxidation. Immune-stimulated macrophages produce nitric oxide, and inhibition of this production reduces the microbiocidal and tumouricidal activities of macrophages. In 1992 Science voted nitric oxide as `Molecule of the Year'; in reviewing the chemistry and biology of NO·, its antibacterial properties were linked to the long-established practice of curing meat with nitrite. Free radicals seem to be cropping up everywhere!

Although the diverse biological roles of nitric oxide is currently the subject of intense activity, only quite recently have the underlying chemical mechanisms been accessible, through the powerful techniques of flash photolysis, its cousin, pulse radiolysis, and electron spin resonance spectroscopy. Readers may be familiar with the brown fumes of nitrogen dioxide in chemical waste stacks or urban smog. This colour reflects reaction of nitric oxide with oxygen, an unusual three-body collision in a formal sense:

but actually proceeding through intermediate steps involving molecules such as N₂O₄ and N₂O₃. The lifetime of nitric oxide, and therefore any biological activity, depends not only on the existence of other free radicals (e.g. superoxide), but also on oxygen tension, and on other reactions which are also important in the photochemical smog characteristic of the Los Angeles basin and in the chemistry of cigarette smoke. Thus free-radical kinetics are also important in the action of environmental carcinogens. Only with the measurement of reaction rates - the quantitation of chemical kinetics - can we make sense of the biology of nitric oxide and superoxide.

The stimulation of superoxide formation by redox-cycling bioreductive drugs, through reactions (10) and (11), could lead to vasoconstriction by reducing the steady-state levels of nitric oxide through equilibrium (12). Superoxide has thus been termed `endothelium derived contracting factor'. High doses of nitroimidazoles, which should stimulate superoxide production by redox cycling, do indeed induce a reduction of blood flow in mice. Although clinical doses of these compounds are lower than those associated with measurable effects in mice, the consequences of increased oxidative stress following administration of `redox cyclers' have been neglected. There is some evidence that free iron levels may influence the unwanted side effects of such drugs. Clearly, more effort needs to be directed towards improving our understanding of the molecular toxicology of radiosensitizers and bioreductive drugs.

Bleach kills all known germs
At least, one well-known brand claims to do. Activated macrophages and neutrophils (white blood cells) seem to rely on free radicals to make bleach (hypochlorous acid) and then to use it to generate hydroxyl radicals. Free radicals are everywhere.

As shown above in reaction (3), superoxide radicals combine to give hydrogen peroxide. An enzyme (myeloperoxidase) is present in macrophages, which catalyses the reaction between hydrogen peroxide and chloride ion (Cl^-, a constituent of saline and of the cytosol):

The reaction product, HOCl or hypochlorous acid, is the protonated form of the active constituent in bleach, produced when chlorine is bubbled through alkali:

Since the pK of equilibrium (19) is a similar value to physiological pH, hypochlorous acid exits in roughly equal amounts in the undissociated (HOCl) and dissociated (OCl^-) forms at pH 7-8.

Highly reactive hydroxyl radicals can be formed from HOCl/OCl^- on reaction with reductants which are one-electron donors. Important examples include superoxide radicals and ferrous iron:

Pulse radiolysis was used by Benon Bielski and colleagues several years ago to measure the rate of reaction (20). However, it was only this year that the Gray Laboratory Molecular Mechanisms Group demonstrated the existence of free hydroxyl radicals in the reaction. When hydroxyl radicals react with a benzene ring, such as that in benzoate (benzene substituted with a -CO₂^-   group), a mixture of substituted hydroxybenzenes are formed. The hydroxyl group can be substituted at three different positions in the ring relative to the position of the CO₂ substituent. Thus a mixture of products is formed in a fixed ratio. We obtained this `fingerprint' of the distribution of the three isomeric hydroxylated products from benzoate, and compared the products from the superoxide/hypochlorous acid reaction (20) with that from truly free ·OH produced radiolytically.

Bleach (hypochlorous acid) is chemically reactive in its own right, and only by a detailed evaluation of its chemical reactivity and consideration of the spatial constraints of its generation in biology can we understand why, or whether, a damaging reactant such as bleach is made even more reactive by generating free hydroxyl radicals. Chemical kinetics has raised its head again in this context. If we compare reaction (21) with the Fenton reaction (5) we see that the reactions are analogous, with H₂O₂ in (5) being replaced by HOCl in (21). Both reactions `convert' ferrous iron to hydroxyl radicals, and involve dissociative electron attachment to oxidants, H₂O₂ and HOCl respectively. However, work this year at the Gray Laboratory has shown that reaction (21) is about a thousand times faster than (5), at least with a simple chemical form of ferrous iron in one model system. Whether enzymes such as myeloperoxidase evolved in order to enhance the formation of hydroxyl radicals from iron and hydrogen peroxide, in effect replacing (5) by [catalyzed] (17) and (21), is a question only now addressable because of these chemical-kinetic measurements.

Newton was right about chemistry, too
Newton's Laws of Motion describes a reaction being equal and opposite to an action. In chemistry, too, many reactions are equilibria or can be reversed. This is especially true with cellular oxidative stress. Oxidative damage from cellular free radicals, as with radiation damage, can be repaired by natural antioxidants. The most important are vitamin C (ascorbate), vitamin E (à-tocopherol), and glutathione (a thiol, GSH). These can donate either an electron or a hydrogen atom to cellular molecules oxidized by free radicals, including those generated by ionizing radiation. They can thus influence damage to all cellular constituents, including DNA, proteins and lipids (membranes). Lipid peroxidation can be effectively inhibited. In 1979, Robin Willson and colleagues at Brunel University (working with CRC support) showed using pulse radiolysis that after vitamin E (vit-E(OH)) has repaired oxidative damage by donating its hydrogen atom, it can itself be restored to carry out more radical repair reactions by reaction with vitamin C (ascorbate, AH^- ):

When thiols such as GSH repair damage, by donating an hydrogen atom, a sulphur-centred thiyl radical, GS· is formed. One of the major contributions to thiol biochemistry of Gerald Adams and his colleagues (whilst working in the Gray Laboratory in 1967) was the observation by pulse radiolysis of reaction of GS· with glutathione itself, in a reaction which produces the glutathione disulphide radical-anion. The latter reacts extremely rapidly with oxygen and produces superoxide radicals:

Hence radical repair by glutathione in the presence of oxygen itself inevitably produces some cellular oxidative stress. Again, this may be the cell's choice of the lesser of two evils in many circumstances.

There is much current speculation as to the role of antioxidants in diverse diseases and the ageing process itself. The levels of oxidized DNA bases excreted in the urine are being correlated with lifespan. Other correlations involve the extent of mitochondrial generation of superoxide radicals. Whatever the involvement of free-radical induced cellular oxidative stress in the many diseases currently being investigated from this viewpoint, it would be wise to eat your greens, and certainly not to smoke. Smokers seem to have a decreased antioxidant status compared to their saner peers, perhaps from taking in 10¹⁷ free radicals per puff. These are mainly nitric oxide, taken in with 500 ppm isoprene to generate NO₂·, which starts off a nasty chain reaction just like photochemical smog in their lungs: ugh.

The fuller picture
A diagram outlining the chemistry of cellular oxidative stress - how nature says NO - will be found in the Molecular Mechanisms report. Interested readers will find fuller discussion in the volumes listed below: a suggested starting point is the book by Halliwell and Gutteridge.

Conclusions
The involvement of free radicals in cellular oxidative stress, leading to cytotoxicity, is a rapidly evolving field. It is closely connected with the role of free radicals as cellular messengers to control non- cytotoxic physiological responses. Some readers may find the chemistry in this article complicated and difficult to understand. Believe me, that's the easy bit. To place these reactions in a biological context requires much more effort than to measure a rate constant, even for reactions which are over in microseconds. Only by chemistry and biology working hand-in-hand in multidisciplinary institutes can we achieve a full understanding.

Cochrane, C. G. and Gimbrone, M. A., Jr. Eds. (1992) Biological Oxidants: Generation and Injurious Consequences Academic Press, San Diego.

Davies, K. J. A. Eds. (1991). Oxidative Damage & Repair. Chemical, Biological and Medical Aspects Pergamon, Oxford.

Halliwell, B. and Aruoma, O. I. Eds. (1993). DNA and Free Radicals (Ellis Horwood, Chichester).

Halliwell, B. and Gutteridge, J. M. C. (1989) Free Radicals in Biology and Medicine (Clarendon Press, Oxford).

Poli, G., Albano, E. and Diazani, M. U. Eds. (1993). Free Radicals: from Basic Science to Medicine (Birkhäuser Verlag, Basel).

Sies, H. Ed. (1985). Oxidative Stress (Academic Press, London). Sies, H. Ed. (1991). Oxidative Stress: Oxidants and Antioxidants (Academic Press, London).

Spatz, L. and Bloom, A. D. Eds. (1992). Biological Consequences of Oxidative Stress. Implications for Cardiovascular Disease and Carcinogenesis, (Oxford University Press, New York).

Peter Wardman