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Chapter II
Oxygen Radicals and Related Species
Ohara Augusto and Sayuri Miyamoto
Departamento de Bioquímica, Instituto de Química, Universidade de São Paulo, Caixa
Postal 26077, 05513-970, São Paulo, SP, Brazil
1. Introduction1: Oxygen and its Metabolites
Imprinted the Evolution of Life
As discussed in Chapter 1, around three billion years ago, life on Earth consisted of
anaerobic microbes subsisting on the energy provided by cycles of electron transfer between
prevalent populations of electron donors and acceptors. Potential donors were molecular
hydrogen (H2), hydrogen sulfide (H2S) and methane (CH4), whose electrons were transferred
to acceptors such as carbon dioxide (CO2) and, to a lesser extent, sulfate (SO42-). The biggest
electron-donor pool, water (H2O), remained biologically inaccessible until the first
photosynthetic organisms evolved between ~3.2 and 2.4 billion years ago. Employing the
energy of the sun and H2O as the reductant to fuel metabolism, these ancient organisms
produced molecular oxygen (O2) as a waste product. Over a relatively short period of
geologic time, around 100 million years, O2 atmospheric levels increased, indicating that the
energy-transducing process of photosynthesis was replacing the old metabolic networks. As a
consequence, the life forms that evolved in anaerobic conditions had to adapt to O2, hide or
become extinct. With O2 overwhelming the Earth’s reducing atmosphere, these primitive
microbes became restricted to hypoxic or anaerobic environments. However, some of them
were able to escape and adapt to the oxidizing atmosphere. They either modified parts of their
metabolism, or formed symbiotic associations that permitted the coupling of nutrient
oxidation to reduction of O2 back to H2O. This respiratory pathway was far more efficient in
extracting energy from nutrients, but it came with costs. The triplet ground state of O2 is
prone to one-electron transfers, which yield species that are toxic to life (see section 6.3). This
1
Due to space constraints, we cited only a small fraction of the original investigations that contributed to the
current understanding of the multiple roles of radicals/oxidants in Biology. Whenever possible, we cited
reviews that referred to the original papers.
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Ohara Augusto and Sayuri Miyamoto
pushed the development of mechanisms to protect the genetic investment in the metabolic
pathway that drive the “H2O-H2O” cycle upon which most life on Earth would come to
depend. The protective pathways evolved further, driving the evolution of single-cell
eukaryotes, from which all current plants and animals on Earth descend [1, 2].
Adaptation to O2, however, went beyond the creation of redox energy couples that
allowed complex life forms to evolve. As a recent meta-metabolomic analysis demonstrated,
it also pushed an evolutionary explosion of random, alternative and novel metabolic networks
yielding a wide variety of gene products that increased fitness of the organisms. Although we
have a long way to go to understand how life evolved, it is clear that O2 and its derived
metabolites imprinted the evolution of complex life forms [1, 2]. Thus, the participation of
oxygen-derived metabolites in cell regulatory and signaling pathways could be anticipated
[2]. Nevertheless, until recently, oxygen-derived metabolites were mostly considered for their
involvement in cell damaging mechanisms [3, 4].
In this Chapter, the chemical properties of O2 that cause its propensity to produce free
radicals and oxidants as metabolites will be presented. These species, known by the general
term reactive oxygen species (ROS) will be historically contextualized. The discoveries
pointing to their roles as mediators of physiological and pathophysiological circuits will be
summarized. Finally, the reactivity of specific ROS towards biomolecules will be presented in
the context of health and disease - topics that will be extended in the other Chapters of this
book.
Figure 1. Electronic distribution of the electrons of two oxygen atoms (8O) in atomic orbitals, which combine
in molecular orbitals to form molecular oxygen (O2) (left side). O2 contains two unpaired electrons. This is
because in distributing its electrons in molecular orbitals, the last two electrons have two molecular orbitals
of the same energy to occupy (2*) and each electron occupies one (Hund´s rule). Because the presence of
two electrons in antibonding orbitals (2*) energetically cancels out one of the bonding orbitals (2), the two
oxygen atoms in O2 are bound by two covalent bonds (3 bond-occupied orbitals minus 2 antibound semioccupied orbitals). B. Electronic distribution of the electrons of one nitrogen atom (7N) and one oxygen atom
(8O) in atomic orbitals which combine in molecular orbitals to form nitric oxide (NO●) (right side).
Distribution of the 17 electrons of NO● in molecular orbitals leaves an unpaired electron in the 2* orbital
making it a free radical in the ground state. Nitrogen and oxygen in NO● are bound by 2.5 covalent bonds (3
bond-occupied orbitals minus 0.5 antibond orbitals).
Oxygen Radicals and Related Species
3
2. Molecular Oxygen, a Sluggish Oxidant Producer
of Reactive Species
Long before the formation of molecules and ions was conceptualized by quantum
mechanics, the discoverers of molecular oxygen in the late 18th century - Priestley, Scheele
and Lavoisier - reported its beneficial and toxic effects on living organisms. These opposing
effects result from the properties of molecular oxygen, which is formed by the combination of
two oxygen atoms in covalent bonds (O2). The characteristics of a covalent bond, such as
strength, length and direction, depend on the occupied molecular orbitals. O2 is bound by two
covalent bonds (Figure 1), and requires 402 kJ/mol to break into two oxygen atoms (O). This
is roughly the amount of energy required to bring 1 liter of H2O to boiling, indicating that the
bond strength between the two O atoms in O2 is strong. In other words, O2 is a considerably
stable molecule, as attested by our daily experience. Most covalent bonded molecules are
equally stable, possessing energy bonds from 150 to 950 kJ/mol.
However, O2 has its peculiarities. Although stable, a spark triggers its reaction with fuels
(combustion) liberating energy to move our cars. Likewise, the oxidation of the foods we eat
(respiration) produces energy to sustain our lives. Indeed, most reactions of O2 are slow but,
once initiated or catalyzed, liberate a considerable amount of energy. In other words,
reactions of molecular oxygen are favored by thermodynamics but not by kinetics. This is
because O2 has two unpaired electrons in the ground state, which makes it a triplet molecule
(Figure 1A). In contrast, most organic molecules possess all electrons paired, being singlet in
the ground state. To react with them, O2 has to receive a pair of electrons, but this requires
spin inversion, an event that is prohibited by the spin conservation rule (Figure 2). According
to its structure, O2 tends to receive electrons by one-electron steps, reacting rapidly only with
species capable of one-electron transfer. Examples of such species are biomolecules that
produce stable free radicals, such as flavin enzymes and coenzymes, and species that contain
unpaired electrons, such as other free radicals and transition metal ions (Figure 2). Thus, it is
not surprising that iron artifacts are rapidly turned into scrap and that most enzymes that
catalyze biological oxidations contain transition metal ions and/or flavin coenzymes in their
active sites. O2, free radicals and transition metal ions are closely related. In fact, O2 and
transition metal ions are also free radicals because they contain unpaired electrons.
Nevertheless, ‘free radical’ is more frequently defined as a species (molecule, cation or
anion) that contains one unpaired electron. Only in these cases is the unpaired electron
denoted by a superscript dot to the right preceding any charge in the radical formula [5]. The
term free radical is historical because organic radicals, which are used to refer to the chemical
groups in a molecule, are bound as opposed to free. Currently, both terms free radical and
radical are used in the literature and understood by their contexts. A classical free radical in
the ground state is nitric oxide (NO) (Figure 1B), whose role in signal transduction pathways
was established in the last decade of the 20th century [3, 4, 6].
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Figure 2. Schematic representation of the main reactions of O2. Due to its unpaired electrons, O2 does not
react with most organic molecules because they have paired electrons. Such reactions are extremely slow
because they would require spin inversion which is prohibited by the spin conservation rule. O2 reacts rapidly
with molecules that produce stable radicals by one-electron transfer (BM●+) or have unpaired electrons, such
as free radicals (R●) and transition metal ions (Mn+). If O2 receives energy (from light, excited molecules,
etc.), it is excited to singlet oxygens (1gO2 or 1g+O2) that are more reactive towards biomolecules because
there is no spin restriction.
3. Free Radicals in Biology: An Historical Account
In the first half of the 20th century, organic radicals were recognized as intermediates of
chemical reactions, and many of their properties were studied and explored in polymerization
processes. Most investigators considered radicals irrelevant to biology, except for some
pioneers. For instance, the known propensity of O2 to react by one-electron steps (Figure 2)
led some investigators, such as Otto Warburg and Leonor Michaelis, to propose that during
respiration O2 would produce one-electron intermediates, such as superoxide (O2-) and the
hydroxyl radical (HO) (Reaction 1).
O2 + 4H+ + 4e
e
O2
e /2H+
O2∙
2 H2O
e /H+ H2O e  /H+
H2O2
HO∙
H2O
(Reaction 1)
Evidence that O2 produced radicals in animals was first provided by Rebeca Gerschman
and co-workers in 1954 while studying the toxic effects of high O2 pressures and X-ray
irradiation in mice [7]. In the resulting paper, they concluded: “it would appear that
irradiation and O2 poisoning produce some of their lethal effects through at least one common
mechanism, possibly that of the formation of free radicals”. In addition, they showed that
Oxygen Radicals and Related Species
5
protectors against irradiation, such as glutathione and ethanol, protected against oxygen
toxicity, or as currently defined, they acted as antioxidants. Another pioneer was Denham
Harman, who proposed in 1954 that aging was a consequence of free radical attack on
biomolecules (see Chapter 35).
These views remained largely ignored up to 1970, while the properties of free radicals
were extensively studied by radiation chemists. The employment of biochemical approaches
to explore fundamental biological questions, such as the molecular basis of life, provided a
turning point. In this background, McCord and Fridovich discovered the enzyme superoxide
dismutase (SOD) in 1969 [8] (see also Chapter 17). They showed for the first time that
mammals possess an enzyme whose function is to catalyze the dismutation of O2●- (k= 1.6 x
109 M-1 s-1), which also occurs spontaneously at a slower rate (k= 5 x 105 M-1 s-1) (Reaction
2). Dismutation is a typical reaction of most free radicals; two molecules of the same radical
(such as O2●-) react in an electron transfer process to produce oxidized (O2) and reduced
(H2O2) products.
O2∙ + O2∙ + 2 H+
H2O2 + O2
(Reaction 2)
The fact that evolution preserved an enzyme whose function is to dismutate a free radical
indicated that radicals were constantly produced during normal metabolic processes. This
fundamental discovery inaugurated a new area of research, “free radicals in biology”, which
continues to expand.
The discovery of SOD provided evidence that cells generate O2●-. Soon thereafter, it
became clear that during biological processes cells also produce additional related
intermediates, generically termed reactive oxygen species (ROS), such as H2O2, HO● and
hypochlorous acid (HOCl). ROS were shown to be produced during mitochondrial
respiration, phagocyte-mediated killing of pathogens, and xenobiotic metabolism. These
investigations brought new concepts into pathophysiology. One was the importance of
reactions catalyzed by transition metal ions (Fenton chemistry; Chapter 5) to produce HO●, an
extremely potent oxidant, from the less reactive O2●- (Reactions 3 and 4).
Fe 3+ + O2∙
H2O2 + Fe2+
Fe2+ + O2
Fe3+ + HO∙ + HO
(Reactions 3 and 4)
Another was the notion that if not eliminated by the antioxidant defenses, ROS would
attack DNA, lipids and proteins (Chapters 6-9) causing cell and tissue injury. The most
influential one was the concept of oxidative stress defined as an imbalance between free
radicals/oxidants and antioxidants in favor of the former [9]. As a consequence, radicals and
oxidants became associated with most human diseases and many intervention studies were
designed to examine the effects of antioxidant vitamins on diseases, particularly
cardiovascular and neurological.
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Ohara Augusto and Sayuri Miyamoto
By the last decade of the 20th century, however, most data collected from intervention
studies with antioxidants were inconclusive, suggesting that the classical view of oxidative
stress required revision [4, 6]. This was reinforced by the discovery of NO● (Figure 1), a
gaseous free radical, as a major autocrine and paracrine mediator of vascular relaxation,
immune regulation and many other physiological effects. In mammals, NO● is mainly
produced from arginine oxidation, catalyzed by a family of enzymes, the nitric oxide
synthases (NOS) (Reaction 5) [10, 11].
Arginine + NADPH + O2
NOS
NO∙ + Citruline + NADP+ + H2O
(Reaction 5)
Soon thereafter, it was demonstrated that O2●- reacts rapidly with NO● (Reaction 6) and
not only regulates NO● bioavailability but also leads to potent oxidants such as peroxynitrite
(ONOO-/ONOOH)†, HO●, nitrogen dioxide (NO2●) and carbonate radical (CO3●-) (Reactions
7 and 8) [12-14] (see also Chapter 3).
Figure 3. Schematic representation of the roles of radicals and oxidants as mediators of physiological and
pathophysiological circuits. Present evidence indicates that free radicals and oxidants are constantly produced
in vivo from exogenous and endogenous sources and are, directly or indirectly, derived from O2. The main
endogenous sources of free radicals/oxidants are: the mitochondrial electron transport chain and enzymatic
reactions catalyzed by nitric oxide synthases (NOS), NADPH oxidases (NOX) (from phagocytes and other
cell types), xanthine oxidase (XO) and hemeperoxidase enzymes, such as myeloperoxidase (MPO).
Organisms evolved enzymatic antioxidant defenses and the capability to use antioxidants from the diet to
control the physiological levels of free radicals/oxidants. High levels of some oxidants may cause
dysfunction because they oxidize biomolecules that lead to cell and tissue injury if not repaired by the
evolved repair systems. Low levels of certain free radicals/oxidants also compromise physiological functions
that evolved to depend on them. On the other hand, transient and small increases of some oxidants trigger
redox-sensitive signaling pathways.
†
The term peroxynitrite refers to the sum of peroxynitrite anion (ONOO-, oxoperoxonitrate (-1)) and
peroxynitrous acid (ONOOH, hydrogen oxoperoxonitrate) unless otherwise specified. Other
abbreviations are defined in the text.
Oxygen Radicals and Related Species
7
k= 1.9 x 109 M-1 s-1
O2∙ + NO∙
pKa= 6.4
ONOO + H+
ONOO
k= 0,17 s-1 0.7 NO  + 0.7 H+
3
ONOOH
k= 2.6 x 104 M-1 s-1
ONOO
+ CO2
[ONOOCO2]
0.3 HO∙ + 0.3 NO2∙
0.65 NO3 + 0.65 CO2
0.35 NO2∙ + 0.35 CO3∙
(Reactions 6-8)
Furthermore, in the same period, several lines of evidence indicated that H2O2 can act as
a second messenger for receptor agonists, such as growth factor and hormones, prompting the
development of the concept of redox signaling. This involves cellular signal transduction
networks in which the integrative element is a series of interconnected electron transfer
reactions. By exerting second messenger effects, oxidants can regulate major cellular
pathways [15-17]. Such function is likely rooted in the oxygen-dependent evolution of
complex life forms. Our current knowledge about redox signaling mechanisms is still in its
infancy.
For decades, free radical research focused on understanding the formation of reactive
species in vivo, and elucidating how proteins, lipids and DNA are damaged by them, resulting
in cellular injury and disease. The complete picture is more complex, though. Exceeding
levels of reactive species may cause dysfunction, but low levels of certain oxidants also
compromise physiological functions that evolved to depend on them, such as microbicidal
activity, proliferative responses and vasodilation. Moreover, transient and small increases in
some oxidants trigger redox-sensitive signaling pathways. Thus, radicals and oxidants are
presently considered to control signaling circuits involved in physiological and pathological
responses (Figure 3) [4, 6]. Unraveling these interrelated processes requires a better
understanding of cellular oxidative mechanisms, including the identification of the involved
oxidants, the pathways regulating their generation, and their molecular targets. Further
advances in the field will depend on an interdisciplinary effort that combines rigorous
chemical thinking and tools with relevant biological data and insights [16, 17].
4. Oxygen Radicals and Related Species:
General Aspects of the Reactivity of Oneand Two-electron Oxidants
Although discrimination between different biologically relevant oxidants became critical,
the difficulties in detecting these short-lived species in biological media stimulate the use of
general terms, such as ROS, in the literature. As previously noted by Christine Winterbourn
[16, 17], terms such as ROS and antioxidants are appropriate to refer to general classes of
compounds but are counterproductive for understanding mechanisms. This is because not all
antioxidants act by the same mechanism, and the species encompassed by ROS have widely
different reactivities. The same is true for another frequently used general term, RNS (reactive
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Ohara Augusto and Sayuri Miyamoto
nitrogen species). RNS comprise reactive metabolites derived from NO●, such as
peroxynitrite and NO2● (Chapter 3). However, RNS and ROS are closely interrelated (see, for
instance, Reactions 6-8), and only one general term may be more appropriate. Indeed, radicals
typically react rapidly with other species containing unpaired electrons (Figure 2), making it
often difficult to establish which oxidant is involved in a particular physiological process. To
make progress in this direction, it is necessary to recognize the widely different characteristics
of the species encompassed by the term ROS (and RNS). Some are radicals and undergo
one-electron reactions, whereas others are non-radicals and promote two-electron
oxidations. Some are strong and others weak oxidants, and the second-order rate constant of
their reactions with biomolecules vary considerably [16, 17]. For instance, the second order
rate constant of some physiologically relevant oxidants with glutathione (GSH) varies from
undetectable to diffusion-controlled (k ~ 109-1010 M-1 s-1 in aqueous media) (Table 1). The
tripeptide GSH (-L-glutamyl-L-cysteinylglycine), which is present at millimolar
concentrations in most cell types, is a major intracellular reductant because of its cysteine
thiol group (-SH) (Chapter 13).
In this book, oxygen radicals and related species (this Chapter), nitrogen radicals and
related species (Chapter 3) and sulfur radicals and related species (Chapter 4) are discussed
separately. Nevertheless, it is useful to stress some general aspects concerning the reactivity
of biologically relevant oxidants. The oxidizing strength of a radical is given by its oneelectron reduction potential (Eo), with higher values corresponding to more potent oxidants.
This is usually reflected in the rates at which radicals react because of the low activation
energy of radical reactions (Table 1). Thus, it is not surprising that HO● reacts with most
biomolecules with second-order rate constants close to the diffusion limit. That is, the
reaction occurs as fast as the reagents encounter each other. As one-electron oxidants,
radicals favor radical chain reactions as occurs during lipid peroxidation (Chapter 7). A
remarkable exception is NO●, which rapidly interacts with other radicals interrupting radical
chain reactions (Reaction 9; see also Chapter 8).
k~ 109 M-1 s-1
NO∙ + ROO∙
NOOOR
(Reaction 9)
Although low NO● levels inhibit lipid peroxidation, the radical can produce potent
oxidants by reacting with O2●- (Reaction 5) and O2 (Reaction 10).
k= 2  106 M-2 s-1
2
NO∙
+ O2
2 NO2∙
(Reaction 10)
In the case of two-electron oxidants, their reduction potential also determines their
oxidizing strength; however, kinetics determines their reactivity because of the high
activation energy involved in these oxidations. For instance, H2O2 has a higher reduction
potential than peroxynitrous acid (ONOOH) but reactions of H2O2 have higher activation
energies and slower rates (Table 1).
Oxygen Radicals and Related Species
9
Table 1. Relative reactivity of selected radical and non-radical oxidants
Oxidant
Radicals (one electron)a
NO●/3NO
RS●/RS (Cys)
O2●, 2H+/H2O2
HO2●, H+/H2O2
ROO●, H+/ROOH
NO2●/NO2
RO●, H+/ROH
CO3●, H+/HCO3
O3●, 2H+/H2O, O2
HO●, H+/H2O
Non-radicals (two electron)b
ONOOH, H+/NO2, H2O
HOCl, H+/Cl, H2O
H2O2, 2H+/2 H2O
Reduction potential
(E’, V)
kGSH (M-1 s-1)c
-0.80
0.92
0.94
1.06
1.00
1.04
1.60
1.78
1.80
2.31
non detectable
8.0 x 108
10 to 103
n.d.
n.d.
3.0 x 107
n.d.
4.6 x 107
7.0 x 107
1.0 x 1010
1.40
1.28
1.77
6.6 x 102
3.0 x 107
0.9
a
Data for one-electron reduction potential collected from [74].
Data for two-electron reduction potential collected from [17, 36].
c
Second-order rate constants for GSH collected from [17, 22].
n.d., not determined.
b
The kinetics of the reactions of radicals and oxidants with biotargets contribute to the
understanding of oxidant actions under physiological conditions. The relevance of the target
depends on its local concentration ([BM]) and on the second-order rate constant of its reaction
with the oxidant (k). Indeed, the product (k x [BM]= k´(s-1)) allows the ranking of the targets
of an oxidant in homogenous media. For instance, the main target of intracellular
peroxynitrite is likely to be CO2 because it reacts rapidly with peroxynitrite (Reaction 8) and
has a high intracellular concentration (~1.3 mM) due to the bicarbonate buffer. The value of
k´ for CO2 (k´= 2.6 x 104 x 1.3 x10-3 ~ 34 s-1) is unmatched by that of intracellular GSH (~5
mM) (k´= 6.6 x102 x 5 x10-3 ~ 3.3 s-1) (Table 1) and other possible targets, except for some
heme and thiol proteins. Among the latter, peroxiredoxins (Prx), which are abundant and react
rapidly with peroxynitrite (k= 105 -107 M-1 s-1) (Reaction 11), deserve special attention [1620] (see also Chapter 22).
k~ 105-107 M-1 s-1
ONOOH + Prx-S
PrxSOH + NO2
(Reaction 11)
At 5 µM concentration, a Prx whose k= 1 x107 M-1 s-1 (k´= 50 s-1) will compete almost
equally with CO2 for peroxynitrite. It is noteworthy that Prx also react rapidly with H2O2 (k=
105 -107 M-1 s-1). The high second-order rate constant of the reactions of these enzymes with
peroxides is one of the reasons why they are being examined as mediators in redox signaling
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Ohara Augusto and Sayuri Miyamoto
mechanisms [16-20]. It is also important to point out that the k´value permits an estimation
of the diffusion distance of selected oxidants under specific circumstances [21]. As an
example, Figure 4 shows the putative diffusion distance of peroxynitrite in the presence of
physiological concentrations of GSH, CO2 and a Prx. It becomes evident that if the only
peroxynitrite target in a cell were GSH (5 mM), the oxidant would diffuse away from one cell
to others and thereby oxidize distant targets by two-electron mechanisms. In the presence of
both CO2 and GSH, the radicals resulting from peroxynitrite would diffuse for only a few
hundred nanometers, and GSH would be oxidized by one-electron mechanisms to the
glutathionyl radical (GS●) (Figure 4). These examples show the importance of considering
kinetics in planning and interpreting experiments related to oxidant action under
physiological conditions. There are databases (NIST Standard Reference DataBase [22]) and
several reviews articles that collect second-order rate constants for reactions of radicals and
oxidants (see, for instance, [13, 14, 16-18]). It is always useful to consult them, although
other factors that affect reactions in physiological media, such as compartmentalization and
media heterogeneity, remain more difficult to assess.
Figure 4. Estimated diffusion distance of peroxynitrite in the presence of physiological concentrations of
GSH (5 mM), CO2 (1.3 mM) and a peroxiredoxin (Prx) (5µM) in a generic tissue composed of cells with a 20
µm diameter. The scheme shows peroxynitrite migration from the center of the circle over different distances
in the presence of the specified targets. Diffusion distances (l) for a tenfold decrease in the oxidant
concentration are represented as the radius (µm) of a circle and were calculated from the expression
l=2.3(D/k[BM])1/2, where D is the oxidant diffusion coefficient, k is the second-order rate constant value and
[BM] is the concentration of the target [17, 21]. D values for peroxynitrite are unknown and a value of 1000
µm2/s was used for the calculation. For comparison, the diffusion distances of the radicals produced from
peroxynitrite reaction with CO2, nitrogen dioxide (NO2) (188 nm) and carbonate (CO3) radical (152 nm) in
the presence of GSH, are also shown.
Oxygen Radicals and Related Species
11
5. Chemistry of Biologically Relevant
Oxygen-Derived Radicals
The biologically relevant oxygen-derived radicals include O2●- and its conjugated acid
(HO2●), HO●, CO3●-, peroxyl (ROO●) and alkoxyl (RO●) radicals. The sources, properties and
main reactions of these species in physiological environments are discussed below.
5.1. Superoxide (O2●-) and Hydroperoxyl (HO2●-) Radical
Superoxide (O2●-) is formed by the one-electron reduction of O2 (Reaction 1). In
biological systems it is produced by enzymatic reactions catalyzed by oxidases, such as
NADPH oxidases (NOX) and xanthine oxidase (XO), and non-enzymatically by redox active
compounds, such as the semi-ubiquinones of the mitochondrial electron transport chain. O2●
can act as oxidant or reductant, and the dismutation reaction is an example of this double
action (Reaction 2). The second-order rate constant of the dismutation reaction is higher at
acidic pH due to the increase in the concentration of the hydroperoxyl/perhydroxyl radical
(HO2●) (pKa = 4.8) in equilibrium with O2●- (Reaction 12). The determined second order rate
constant value for the spontaneous dismutation varies from  102 M-1 s-1 at pH 11 to  5 x 105
M-1 s-1 at pH 7. As noted above, the dismutation reaction catalyzed by SOD is much faster
(k= 1.6 x 109 M-1 s-1).
O2∙ + H+
HO2∙
(Reaction 12)
O2●- reactivity towards a variety of organic and inorganic targets has been studied.
Despite its moderately high reduction potential (0.94 V) (Table 1), its reactivity with nonradical targets is limited [23]. However, O2●-may induce harmful effects by reacting with
radicals, such as NO● (Reaction 6), and species containing transition metal ions. For instance,
O2●- is highly reactive towards iron-sulfur ([Fe-S]) clusters [24] (see also Chapter 31), with
the reaction occurring at near diffusion-limited rates (Reaction 13). O2●- causes one-electron
oxidation of the [4Fe-4S] clusters to form H2O2 and an unstable intermediate that decomposes
losing iron (II). Relevantly, this free iron ion can react with H2O2 and catalyze the production
of HO● by Fenton chemistry (Reactions 3 and 4).
k > 109 M-1 s-1
[4Fe-4S]2+ + O2∙ + 2 H+
H2O2 + [4Fe-4S]3+
(Reaction 13)
[3Fe-4S]+ + Fe2
Radical-radical reactions of O2●- typically occur at near diffusion-controlled rates [25].
Among them, the reaction of O2●- with NO· has received special attention due to the
generation of peroxynitrite (Reaction 6), a strong biological oxidant (Chapter 3). Reactions of
O2●- with radicals formed on aromatic amino acids, such as the tyrosyl radical (Tyr●), have
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Ohara Augusto and Sayuri Miyamoto
also been studied. The second-order rate constant of the reaction with Tyr● is about 3-fold
higher than that of Tyr● dimerization to dityrosine [25]. Upon reaction with O2●- the Tyr●
radical can be repaired back to Tyr or be transformed into Tyr hydroperoxides depending on
the position of the Tyr residue in the polypeptide chain.
Reactions of O2●- with biologically important thiol compounds have been investigated
[26]. Literature data on the second-order rate constant for these reactions vary from 10 to 105
M-1 s-1. This variation has been attributed to the use of inappropriate assays and the
complexity of the chain reactions involving sulfinyl and thiyl radicals [27]. According to
Winterbourn and colleagues, the best estimates are in the range of 30 to 103 M-1 s-1. These
values indicate that biothiols do not react rapidly enough with O2●- to compete with SOD.
Therefore, proteins containing [Fe-S] clusters are likely to be O2●- sensors in redox signaling
(Reaction 13). In agreement, bacteria rely on the SoxR transcription factor, which contains
[Fe-S] clusters, to upregulate resistance genes in response to sublethal O2●- levels.
5.2. Hydroxyl Radical (HO●)
HO● is the strongest oxidant produced in biological systems (E’ =2.31 V) (Table 1). It
reacts very rapidly and indiscriminately with most biological targets present at its site of
formation (see, for comparison, Figure 4). Indeed, the rate constants for its reactions with
most of biomolecules, including, lipids, proteins, carbohydrates and DNA are very close to
the diffusion-controlled limit. HO● can be generated in vivo by four major processes:
ionizing-radiation (e.g. UV, X-rays, -rays); transition metal ion-catalyzed reactions
(Reaction 4); proton-catalyzed decomposition of peroxynitrite (Reaction 6); and
decomposition of ozone (O3) (see section 6.5). Irradiation with high-energy radiations
generates HO● by homolytic fission of water molecules (Reaction 14). HO● produced by this
process is thought to be responsible for DNA damage and tumor development (e.g. skin
cancer), while irradiation can also be employed for targeted killing of tumor cells (e.g.
radiotherapy).
H2O
HO∙ + H∙
(Reaction 14)
Reactions of the HO● can be classified into three main types: hydrogen abstraction,
addition and electron transfer reactions. HO● is extremely fast in abstracting hydrogen atoms
from organic compounds, especially from those that are weakly bound. The reaction produces
H2O and a carbon-centered radical (RC●), which, in the presence of O2 generates peroxyl
radicals (ROO●), such as in the case of carbon-centered lipid radicals. In the absence of O2, a
covalent bond can be formed between two RC● producing a crosslink. An example is the
reaction between two Tyr● to produce dityrosine. Reactions of HO● with aromatic compounds
usually involve the addition of the radical to produce hydroxylated radical adducts. An
important example is the addition of HO● to the guanine moiety of DNA/RNA to produce 8hydroxyguanine and 2,6-diamino-4-hydroxy-5-formamidopyrimidine (Chapter 9). Reactions
of HO● with anions produce radicals by electron transfer mechanisms. For instance, the
Oxygen Radicals and Related Species
13
abundant Cl- and the NO●-metabolite, NO2-, are oxidized to chlorine (Cl●) and NO2●,
respectively (Reactions 15 and 16).
Cl
NO2
+
+
k= 4.3  109 M-1 s-1
HO∙
Cl∙ +
k= 1  1010 M-1 s-1
NO2∙
HO∙
HO
+
HO
(Reaction 15 and 16)
5.3. Carbonate Radical (CO3●-)
Although less oxidizing than HO●, CO3●- (Eo´ =1.78 V) (Table 1) is a very strong oneelectron oxidant. Relatively recent data demonstrated that CO3●- is an electrophilic oxygencentered radical and a strong acid (pKa <0). Consequently, it will be negatively charged in all
physiological environments, including those of acidic pH [13, 28]. The most recognized
biological source of CO3●- is the reaction between peroxynitrite and CO2 (Reaction 8) (Figure
4). CO3●- is also produced by the peroxidase activity of SOD, by the turnover of XO and by
transition metal-ion catalyzed decomposition of peroxymonocarbonate (HCO4-) [13, 28]. The
latter is an oxidant found in equilibrium in aqueous solutions of H2O2 and bicarbonate (HCO3) (Reaction 17), whose potential participation in biological processes has been recently
discussed [13, 28].
H2O2
+ HCO3
HCO4 + H2O
(Reaction 17)
Most CO3●- reactions are oxidations by both electron transfer and hydrogen abstraction
mechanisms to produce radicals from the oxidized targets. Addition reactions of the carbonate
radical to produce stable target adducts are virtually unknown in the literature [13, 28]. As a
charged species, CO3●- is an important oxidizing agent in aqueous environments where
intracellular GSH is a particularly relevant target (Table 1). In addition, CO3●- oxidizes
solvent-exposed amino acid residues in proteins, particularly Tyr, Trp and Cys. Guanine
moieties in nucleotides, nucleosides and nucleic acids are efficiently oxidized by CO3●-.
Antioxidants, such as ascorbic and uric acid, are also rapidly oxidized by CO3●-. Because
CO3●- mostly oxidizes biomolecules to radicals, its effects are difficult to discriminate from
those of other strong one-electron oxidants. This fact contributes to the limited attention that
CO3●- has received so far, despite bicarbonate being the main physiological buffer [13, 28].
5.4. Peroxyl(ROO●) and Alkoxyl (RO●) Radicals
ROO● and RO● are moderately strong oxidants (E’ > 1.0 V) (Table 1). They can be
generated from organic hydroperoxide (ROOH) decomposition induced by heat or radiation
14
Ohara Augusto and Sayuri Miyamoto
and by ROOH reaction with transition metal ions and other oxidants capable of abstracting
hydrogen (Reactions 18 and 19). ROO● are also important intermediates in processes
involving carbon-centered radicals, which react rapidly with O2 (k>109 M-1 s-1). In addition,
biomolecule-derived ROO● and RO● can be generated from the oxidation of lipids, proteins
and nucleic acids.
ROOH + Fe2+
RO∙ + HO + Fe3+
ROOH + Fe3+
ROO∙ + H+ + Fe2+
(Reaction 18 and 19)
The reactivity of ROO● and RO● is influenced by the substituents at the -carbon. An
electron-withdrawing group increases the reactivity (for instance, chloroperoxyl radical,
CCl3OO●), whereas electron-donating groups decreases it (for instance, phenoxyl radicals).
Aromatic ROO● and RO●· tend to be less reactive because of unpaired electron delocalization.
The reactions of ROO● and RO● with biomolecules often involve hydrogen-abstraction,
which is facilitated in compounds containing weakly bound hydrogens. This is the case for
lipids, thiols, and several chain-breaking antioxidants. Lipids are particularly susceptible to
hydrogen abstraction and this reaction is the rate-limiting step in the propagation of lipid
peroxidation chain reactions (Chapter 7). The second-order rate constant value of ROO●mediated hydrogen abstraction from lipids is low (k < 102 M-1 s-1) [29, 30] and increases with
the number of allylic or double-allylic hydrogens. For unsaturated fatty acids the secondorder rate constant value decreases in the sequence 22:6>20:5>20:4>18:2>18:1. Notably, the
second-order rate constant values of RO● reactions (k  106-107 M-1 s-1) [29] are about 4-5
orders of magnitude higher than those of ROO●. Both ROO● and RO● can undergo rapid
monomolecular rearrangements or fragmentations that compete with hydrogen abstraction
reactions. ROO● formed on aromatic rings and those with -carbon linked to hydroxy or
amino groups can decompose to liberate O2●- or HO2●. This type of reaction has been reported
for amino acids, such as lysine. ROO●· can also react with another ROO●· by the Russell
mechanism, generating a ketone, an alcohol and singlet molecular oxygen (1O2) (see section
6.3) [31].
6. Chemistry of Biologically Relevant
Non-radical Oxygen Species
Biologically relevant two-electron oxidants derived from oxygen include H2O2, HOCl
and related species, 1O2, biomolecule-derived hydroperoxides and ozone (O3). These species
are discussed below with regard to main properties and reactions in physiological
environments.
Oxygen Radicals and Related Species
15
6.1. Hydrogen Peroxide (H2O2)
Biological production of H2O2 can occur by chemical and photochemical processes and
by enzymatic reactions catalyzed by several oxidases, such as NOX, XO and monoamine
oxidase (MAO). In addition, H2O2 is produced continuously from O2●- dismutation, either
spontaneous or catalyzed by SOD (Reaction 2). In physiological environments, H2O2 can be
rapidly inactivated to H2O by seleno-, heme- and thiol- peroxidases. For instance, the secondorder rate constant value of the reaction of H2O2 with glutathione peroxidases (GPx),
catalases and Prx are ~108 M-1 s-1, ~106 M-1s-1 and 105 -107 M-1s-1, respectively. Thus, steadystate concentrations of H2O2 in cells and tissues are low, being estimated to be around 10-7 to
10-8 M.
H2O2 is a powerful two electron-oxidant (E = 1.77 V, pKa 11.6) (Table 1). However, its
reactivity toward most of biological molecules is low because of the high activation energy of
these oxidations. Most damaging effects of H2O2 in vivo are considered to be mediated either
by transition metal ions or enzymes, such as heme-peroxidases. These processes generate
secondary species, which are more reactive and include radicals, such as HO● and NO2●, and
non-radical species, such as HOCl and related species. Reaction of H2O2 with reduced
transition metal ions, such as copper (I) and iron (II), leads to the generation of HO●
(Reaction 4). This reaction has been studied for more than a century (Chapter 5). The secondorder rate constant values of Fenton reactions are dependent on the metal-ligand and on the
medium pH, ranging from 102 to 104 M-1 s-1 for iron ions.
In addition to being substrate for several heme-peroxidases, H2O2 can be consumed by
other hemoproteins, such as hemoglobin, myoglobin and cytochrome c, in processes that
oxidize the proteins. In phagocytic cells, myeloperoxidase (MPO) uses H2O2 mainly for the
production of HOCl (Reactions 20 and 21). MPO has also been reported to use H2O2 to
oxidize NO2- to NO2● [32] and Tyr to Tyr● [33]. Thus, MPO can be an efficient mediator of
protein nitration, particularly at acidic pH [34]. MPO is released during phagocytosis and is
thought to play an important role in microbial killing [35]. However, excessive MPOmediated production of HOCl, NO2● and other oxidants can cause host tissue injury,
contributing to the development of several diseases [36, 37].
MPO + H2O2
MPO + Cl
k1
k -1
k2
MPO-I + H2O
MPO + HOCl
(Reaction 20 and 21)
Among the physiological reactions of H2O2, the oxidation of biothiols is receiving
increasing attention in the literature [15, 16]. The second-order rate constant values of the
oxidation of low molecular weight thiols (RSH) by H2O2 are relatively low (1-5 x 102 M-1 s-1).
The value increases for thiols with low pKa [26], indicating that the effective substrate is the
thiolate form (RS-). In fact, it is proposed that a nucleophilic attack of RS- on peroxide
oxygen occurs to produce sulfenic acid (RSOH) (Reaction 22), which subsequently reacts
with a second thiol to produce disulfide (RSSR) (Reaction 23).
16
Ohara Augusto and Sayuri Miyamoto
H2O2 + RS
RSOH + H2O
RSOH + RSH
RSSR + H2O
(Reaction 22 and 23)
It is important to emphasize that some classes of proteins possess Cys residues whose
pKa values are in the range of ~4.0 to ~6.5, which are much lower than that of free Cys
(pKa=8.4) or GSH (pKa=9.2). Examples of these proteins are the Prx, which react extremely
rapid with H2O2 (k~105-107 M-1 s-1). Currently, it is accepted that Cys residues with low pKa
are only one of the factors contributing to the high reactivity of Prx towards peroxides [18]. It
will be important to unravel Prx catalysis at the molecular level because these enzymes are
likely to participate in H2O2-mediated signaling cascades (see also Chapter 22) [16-20].
6.2. Hypochlorous acid (HOCl) and Related Species
HOCl and related species (HOX, X= Cl, Br, I and SCN) are moderately strong twoelectron oxidants (E’HOCl/Cl=1.28 V, HOBr/Br=1.13 V, HOSCN/SCN=0.56V)
generated during inflammatory processes (Chapter 33). They are primarily produced from the
reaction of H2O2 with halide and pseudo-halide ions (Cl, Br, I and SCN ) catalyzed by
MPO and eosinophil peroxidase [36, 37]. In terms of plasma concentration, the most
abundant halide ion is Cl (100-140 mM) followed by Br (20-100 µM), SCN (20-120 µM)
and I (0.1-0.6 µM) [38]. Thus, it is believed that HOCl is the major hypohalous acid
produced during phagocytosis [37, 39]. However, the second-order rate constant value for the
halogenation reaction is about 10-20 fold higher for SCN and Br compared to Cl and
significant amounts of HOSCN and HOBr can also be produced.
At physiological pH, HOCl is in equilibrium with its conjugate base, hypochlorite (OCl-,
pKa 7.59 at 25C) [36] (Reaction 24), and both forms appear to be responsible for oxidation
and/or halogenation reactions. In acidic conditions, HOCl can be in equilibrium with
molecular chlorine (Cl2, pKa 3.3) [40] (Reaction 25). In vitro studies suggest that Cl2 might
be the agent that mediates formation of chlorinated products during phagocytosis [40].
HOCl
pKa = 7.4
HOCl + Cl + H+
H+ + OCl
pKa = 3.3
Cl2 + H2O
(Reaction 24 and 25)
HOCl is reactive towards several biomolecules. Amino (RNH2) and thiol (RSH) groups
of amino acids and peptides are among the most important targets (Table 1). Oxidation of
these groups by HOCl yields unstable chloramines (RNHCl) and sulfenyl chlorides (RSCl),
respectively [41-43]. Both intermediates induce further reactions that lead to an increased
oxidative damage to biomolecules [36]. HOCl is also reactive towards aromatic rings in
amino acids and nucleobases. Reaction of HOCl with Tyr yields 3-chlorotyrosine and 3,5-
Oxygen Radicals and Related Species
17
dichlorotyrosine [44, 45]. With lipids, HOCl adds across carbon-carbon double bonds in fatty
acids and cholesterol yielding chlorohydrins [46]. HOCl also reacts with H2O2/ROOH. The
reaction with H2O2 produces stoichiometric amounts of 1O2 (Reaction 26) [47, 48]. Similarly,
HOCl can react with lipid hydroperoxides to yield 1O2 through a mechanism involving the
generation of ROO● [49].
H2O2 + HOCl
1O
2
+ Cl + H2O + H+
(Reaction 26)
6.3. Singlet Molecular Oxygen (1O2)
O2 refers to the excited states of O2, the 1g and 1g+ state. They have energies of 94.3
kJ/mol and 156.9 kJ/mol above the triplet ground state, respectively. The 1g+ state has an
extremely short lifetime in H2O (~10-11 s), decaying rapidly to the 1g state, which is
considered the biologically relevant form of 1O2. The lifetime of 1O2 is greatly influenced by
the solvent type, being in the range of 1-5 x 10-6 s in H2O and about ten times higher in
deuterium oxide (D2O). Thus, D2O is commonly employed in experiments devised to prove
1
O2 formation. Photosensitization is the most conventional source of 1O2 in biological
systems. Typically, this occurs by photoexcitation (type II reactions) of endogenous
photosensitizers (porphyrins, flavins, quinones, etc.) exposed to UVA [50]. These processes
are especially important in the skin. Excessive production of 1O2 has been associated with
disease states, such as porphyrias that are characterized by porphyrin accumulation in the
skin. 1O2 promotes tumor cell death and this property is exploited in photodynamic therapy.
1
O2 can also be generated by non-photochemical reactions, usually involving inorganic
peroxides (such as H2O2 and peroxynitrite) and organic hydroperoxides (ROOH). For
instance, the generation of 1O2 has been evidenced during phagocytosis [35, 51], lipid
peroxidation [52] and peroxidase turnover [53, 54]. The reaction between HOCl and H2O2
(Reaction 26) is the likely source of 1O2 during phagocytosis. The generation of 1O2 during
lipid peroxidation has been attributed to the combination of ROO● radicals by the Russell
mechanism [31, 55]. It is important to note that tertiary peroxyl radicals are unable to
generate singlet oxygen because the hydrogen- on one of the ROO● is required for the
elimination of O2 in the singlet-excited state [56]. The yield of 1O2 by this mechanism has
been estimated to be approximately 3.9 to 14% [57].
1
O2 is a strong two-electron oxidant that displays considerable reactivity towards
electron-rich organic molecules, including nucleic acids, proteins and lipids [58]. Typically,
1
O2 adds to -bonds by three common mechanisms: addition to alkenes containing allylic
hydrogen by a ene-type reaction yielding hydroperoxides; 1,4-cycloaddition to 1,3-dienes
(Diels-Alder reaction) forming endoperoxides; and 1,2- cycloaddition to electron-rich alkenes
producing dioxetanes [59]. Chemical reactions of 1O2 occur with second-order rate constant
values that are usually lower than 107 M-1 s-1. For instance, the values obtained for the
chemical reaction of 1O2 with unsaturated fatty acids are in the order of 104 M-1 s-1. In
contrast, physical quenching of 1O2 occurs at much faster rates and values near the diffusioncontrolled limit were observed for carotenoids, among which lycopene exhibited the highest
1
18
Ohara Augusto and Sayuri Miyamoto
rate [60]. Reactions of 1O2 with thiols and ascorbate have been also reported. In the case of
thiols, 1O2 reacts preferentially with the thiolate anion yielding several oxidation products
[61]. Apparent second-order rate constant values calculated for low molecular weight thiols,
such as Cys, N-acetylCys and GSH, were in the order of 106 M-1 s-1 [61]. Ascorbate reacts
with 1O2 (k = 3 × 108 M-1 s-1) producing H2O2 and dehydroascorbic acid [62].
6.4. Organic Hydroperoxides from Biomolecules
Several classes of organic hydroperoxides (ROOH) are produced upon oxidation of
biomolecules, including lipids, proteins and DNA [63]. Lipid hydroperoxides can be formed
enzymatically during lipoxygenase, cyclooxygenase, cytochrome P450 and heme-peroxidase
turnover [64, 65]. Moreover, a great number of ROOH are generated non-enzymatically by
the oxidation of biomolecules mediated by radicals and 1O2 [64]. ROOH are relatively stable;
however, they can participate in reactions that decrease or increase their toxicity. Normally,
cells contain enzymes that reduce ROOH to their corresponding alcohols, decreasing their
reactivity and toxicity. The enzymatic reduction of lipid hydroperoxides has been extensively
studied. Three classes of enzymes are known to mediate their two-electron reduction: GPx,
glutathione S-transferases (GST) [64, 66] and Prx [20]. Toxicity is increased when ROOH
are converted to RO or ROO [3]. This can occur especially in the reaction of ROOH with
transition metal ions, hemoproteins and other one-electron oxidants [3].
ROOH can react with reduced and oxidized states of free transition metal ions to form
oxyl radicals. In the first case, RO● is produced whereas in the second case, ROO● is
produced [3, 64]. Oxyl radicals are responsible for propagating the oxidation process as well
as for the generation of other highly reactive products capable of causing modifications in
proteins and nucleic acids. These include electrophilic aldehydes, epoxides, ketones, and
excited species, such as 1O2 and electronically excited carbonyl species [63, 67].
6.5. Ozone (O3)
O3 is present in polluted atmospheres, and inhalation of this toxic triatomic gas can
induce lung injury and inflammation. In biological systems, O3 is thought to be generated
during antibody catalyzed oxidation of H2O to H2O2 [68]. In the proposed mechanism,
antibodies use H2O as an electron source, facilitating its addition to 1O2 to form hydrogen
trioxide (H2O3) as the first intermediate in a cascade of reactions that eventually leads to O3.
The generation of O3 in human tissues has been postulated based on the detection of chemical
signature products, including isatin sulfonic acid and cholesterol secoaldehyde [69].
However, these markers are not specific for O3, casting doubts on whether O3 can be
generated endogenously by cells [70-72].
Reactions of O3 with fatty acids, cholesterol, amino acids and DNA have been
characterized. O3 adds directly to the double bonds in fatty acids, giving rise to Criegee
ozonide (10%) and a hydroxyhydroperoxide intermediate (90%) that decomposes to form
aldehydes and H2O2 (Reaction 27) [70]. With saturated organic targets and inorganic
compounds, O3 reacts with nucleophiles, especially those containing nitrogen or sulfur atoms
Oxygen Radicals and Related Species
19
by oxygen-transfer mechanism to generate a trioxide intermediate. Decomposition of this
intermediate is reported to generate 1O2 and the corresponding oxidized product [73]. In
aqueous solutions, one electron-transfer reactions can also occur, leading to the formation of
ozonide (O3), HO● and O2●- [73].
O3
H2O
+
+
H2O2
(Reaction 27)
7. Conclusions
In its initial four decades, the field of “free radical research in biology” has focused on
understanding the mechanisms by which radicals and oxidants are produced in vivo and how
proteins, lipids and DNA are oxidized by them resulting in cell damage and, eventually,
disease. The current picture is more complex, though. An increasing body of evidence
suggests that apart from being potentially toxic, radicals and oxidants also exert important
(patho)physiological signaling functions (Figure 3). The diverse biological activities of
oxygen radicals and related species are likely rooted in the oxygen-dependent evolution of
complex life forms. Although the chemical properties and reactivities of relevant biological
oxidants are becoming well understood, it remains less clear how these properties translate
into cellular and tissue responses. Further advances in the field will require interdisciplinary
approaches combining chemical reasoning with biological insights.
Acknowledgments
Our laboratories are supported by grants from Fundação de Amparo à Pesquisa do Estado
de São Paulo (FAPESP), Conselho Nacional de Desenvolvimento Científico e Tecnológico
(CNPq) and Coordenação de Aperfeiçoamento de Pessoal de Nível Superior (CAPES). The
authors are members of the INCT de Processos Redox em Biomedicina-Redoxoma
(CNPq/FAPESP/CAPES).
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