An ever-present suspect, the root of all diseases?
Whenever I hear the term reactive oxygen species, I am reminded of an ingredient that seems to turn up in every prescription. In papers and research reports about diseases, they seem to make an appearance almost without fail. They are even said to attack telomeres and contribute to accelerated ageing. But why? And what exactly are reactive oxygen species? I have a vague sense that they are the villains of the story, but I want to understand precisely what makes them harmful—and whether they really do nothing but harm. So I have decided to take a proper, detailed look. Reactive oxygen. Presumably, oxygen that has become reactive. Then perhaps the best place to begin is with a simple question: what does it mean to be reactive?
What does it mean to be reactive?
When a substance is unstable and readily reacts with other substances, we describe it as reactive. Atoms in stable states generally have lower energy levels and tend to be less reactive. But what makes something reactive in the first place? In earlier articles, we looked at the rules governing how electrons are arranged in orbitals within an atom. We also saw that each orbital can accommodate a maximum of two electrons. In fact, electrons generally prefer to exist in pairs. But suppose an electron is lost or taken away for some reason. Its partner is now left alone, and the system becomes unstable. Its former stability and peace are disturbed, and it is almost as though the unpaired electron begins looking around for a way to find a new partner, becoming increasingly restless in the process. The word radical itself brings to mind something wild, unruly, or difficult to control. We can already begin to see why reactive oxygen species might have something to do with electrons that have lost their partners.
What is a radical?
A molecule containing one or more unpaired electrons is called a radical.[1] Electrons generally tend to exist in pairs. To form substances, atoms combine into molecules through various types of chemical bonding. Consider a covalent bond, in which two atoms each contribute an electron to form a shared pair. This helps explain why even atoms with an odd atomic number do not necessarily remain radicals once they form molecules. In fact, when molecules break apart, the electrons in a bond can remain paired. But when a stable covalent bond is broken in a different way, with one electron going to each fragment, two unpaired electrons are produced. Each fragment now has an unpaired electron. This is how radicals can form.
The frightening nature of radicals: chain reactions
To restore its unstable state after losing an electron partner, a radical may seek a new partner by attacking a nearby molecule and taking an electron from it. The radical may then become more stable, as though it has finally found what it was looking for. But there is a problem. The unfortunate molecule that has just lost an electron may itself become a radical. One radical has given rise to another. This newly formed radical may then attack another nearby molecule to obtain an electron, setting off the same process again. Such chain reactions can unfold within an extremely short time. The formation of the first radical is called initiation. The series of reactions in which radicals continue to generate new radicals is called propagation. As radicals repeatedly attack non-radical molecules, the chain reaction continues. Eventually, however, radicals may react with one another to form stable, non-radical products. When this happens, the chain reaction comes to an end, a stage known as termination. Perhaps this ability to set off chain reactions is the most frightening characteristic of radicals. And such reactions really do occur inside our bodies. One example is the oxidation of lipids in cell membranes, which we encountered in an earlier article. If radical chain reactions continue in the same place for a prolonged period, the consequences could be devastating.
Oxidation and reduction (redox)
When we consider what happens when a paired electron is lost and its partner is left unpaired, we inevitably encounter the concepts of oxidation and reduction. In a chemical reaction, oxidation occurs when an atom or molecule loses electrons, combines with oxygen, or loses hydrogen. As the name suggests, oxidation originally referred to reactions involving oxygen, but it was later discovered that oxidation can occur even without oxygen. Perhaps the more general definition is simply the loss of electrons. Reduction is the opposite: gaining electrons, losing oxygen, or gaining hydrogen. Oxidation and reduction commonly occur together. After all, if one substance loses an electron, another must receive it. It follows that reactive oxygen species can arise during oxidation–reduction reactions, or redox reactions, in which electrons are transferred between substances.
Losing electrons: an essential part of energy production
It is remarkable how much significance such a seemingly simple process can have. Through the rather trivial(?) act of losing electrons, we sustain life itself. Let us take a brief look at one representative example. After we obtain nutrients from food, they are carefully broken down, step by step, until substances such as pyruvate are produced. Pyruvate is then slipped from the cytoplasm into the mitochondria, where the repeating series of reactions known as the TCA cycle begins. The electron transport chain, located in the inner mitochondrial membrane, is the system that helps convert the energy carried by NADH and FADH₂ into ATP, a form of energy our bodies can use. In this process, the electrons carried by NADH and FADH₂ are passed from one protein complex to another, travelling through an elaborate series of proteins and reactions. Eventually, they reach oxygen, the final electron acceptor, which combines with hydrogen ions to form water. At this point, I cannot help wondering: if the electrons are ultimately going to oxygen anyway, why not simply deliver them directly? Why go through such a complicated series of steps? There is a reason. As electrons are transferred towards oxygen, energy is released—and a considerable amount of it. What would happen if all that energy were released at once, in a form our bodies could not manage? The thought is rather alarming. Instead, the electrons take a long, carefully arranged detour, allowing the energy to be released little by little, in manageable steps rather than all at once.
But does our body use the energy released during electron transfer directly? Not quite. Much of this energy is used to pump hydrogen ions, or protons (H⁺), from the mitochondrial matrix across the inner membrane into the intermembrane space. As protons accumulate between the inner and outer mitochondrial membranes, their concentration becomes much higher than it is inside the matrix. This difference creates a powerful driving force, allowing protons to rush back into the matrix through ATP synthase. The energy associated with this flow is used to produce ATP, the energy our bodies rely on. The principle is rather like water falling from a height and turning a waterwheel, or water flowing through a dam and driving a turbine. In much the same way, the energy stored in the proton gradient is used to produce ATP. ADP has two phosphate groups, and to become ATP it must acquire one additional inorganic phosphate. The energy required for this process comes from the proton gradient across the inner mitochondrial membrane. I intend to explore cellular respiration, the process through which our bodies produce usable energy, in much greater detail in a separate article. But there is a reason I have spent so much time describing it here. This very process, essential for sustaining life, is also a major source of reactive oxygen species. In other words, their formation is an unavoidable part of the process. Electrons are carefully carried along, almost as though they are precious parcels being delivered to oxygen. Yet some inevitably escape along the way, taking a wrong turn, getting lost, or leaking out before reaching their intended recipient. Delivery accidents happen quite frequently: electrons that were supposed to be delivered to the oxygen waiting at their destination end up going astray somewhere along the way. When these escaped electrons react with oxygen, they can initiate the formation of reactive oxygen species.
How oxygen becomes reactive oxygen species
Now that we have looked at what radicals are and why substances become reactive, let us turn our attention to oxygen itself. First, we will follow the changes through which relatively stable oxygen can give rise to reactive forms. There are several different types of reactive oxygen species, collectively known as ROS. As we will see, these different species are also closely connected to one another.
The oxygen atom
An oxygen atom has an atomic number of 8, meaning that it contains eight electrons. Applying the rules of electron configuration that we explored in earlier articles, based on electron shells and orbitals, we can represent oxygen as shown in the image below. The first shell, or K shell, closest to the nucleus contains two electrons in its s orbital. The second shell, or L shell, contains another two electrons in its s orbital. The remaining four electrons occupy the three p orbitals: two in px, one in py, and one in pz. This gives us a total of eight electrons. Notice that in two of these orbitals, the electrons remain unpaired.
There is one important point to keep in mind here. The electron configuration diagrams used in this article have been simplified to make it easier to understand the structure of oxygen and the changes that occur as electrons are added. In reality, electrons do not travel along fixed paths or remain at specific positions as the diagrams might suggest. Rather than focusing on their exact locations, it is better to pay attention to how the electrons are arranged and how that arrangement changes.
Electrons generally form pairs in stable configurations. In oxygen, however, the two outermost p orbitals, py and pz, each contain a single electron. If we recall the octet rule, according to which atoms tend to achieve a stable arrangement of eight electrons in their outermost shell, we can see that an oxygen atom needs two more electrons to complete its octet.
Triplet oxygen
Two oxygen atoms with this electron configuration form a covalent bond, producing O₂—the triplet oxygen that we breathe.
By sharing electrons in their outermost orbitals, the two oxygen atoms form electron pairs through covalent bonding. Yet, in this simplified picture, each oxygen atom still has an orbital containing an unpaired electron. From the perspective of the octet rule, it may appear that one electron is still missing. We might therefore expect this molecule to be unstable because of those unpaired electrons. Surprisingly, however, oxygen is relatively stable in this state, even though not all of its electrons are paired. Stable enough, in fact, for us to breathe it and live. But what happens when this stable triplet oxygen gains an electron from its surroundings—in other words, when it is reduced?
Superoxide anion radical
If we simplify the diagram by grouping together the orbitals of the L shell, the result looks like the image above. Initially, each oxygen atom appeared to be one electron short. After the molecule gains an electron through reduction, one side now satisfies the octet rule, but an unpaired electron still remains on the other side. This raises an interesting question. Why does oxygen not simply gain two electrons at once? Because triplet oxygen is a species that typically accepts electrons one at a time. The result is a molecule containing an unpaired electron—the very definition of a radical that we encountered earlier. Even the name seems to suggest instability. That remaining unpaired electron makes the molecule reactive, as though it were determined to find a partner somehow. The word anion tells us that the molecule has acquired a negatively charged electron and therefore carries an overall negative charge. I must admit that when studying chemistry, I often find the long, complicated, and rather unfriendly names of chemical substances intimidating. But once we begin to understand the rules behind chemical naming, the names become less daunting and can even tell us something useful about the properties of the substances themselves. I hope to study this subject in greater detail and write about it in a future article. Incidentally, radicals are indicated by a dot in their chemical formulas.
Starting with the oxygen atom, we have now encountered our first radical. This is the superoxide anion radical, commonly called superoxide. Like other radicals, it can react with various substances and give rise to a range of chemical products. What makes it particularly important is its involvement in the formation of other reactive oxygen species, or ROS. ROS are reactive substances containing oxygen. Superoxide is itself a radical, but we should remember that it can also lead to the formation of hydrogen peroxide by gaining another electron and combining with hydrogen ions. Hydrogen peroxide, in turn, can give rise to the hydroxyl radical, one of the most reactive and damaging radicals. This is why superoxide deserves particular attention.
Formation of hydrogen peroxide
1. Formation of peroxide
When the superoxide radical gains another electron from its surroundings, its previously unpaired electron becomes paired. The result is the peroxide ion (O₂²⁻). Peroxides are substances characterised by an O–O bond between two oxygen atoms. Different atoms or groups of atoms can attach to these oxygen atoms, producing various types of peroxides. Hydrogen peroxide (H₂O₂), for example, has one hydrogen atom attached to each end of the O–O structure.
2. Hydrogen peroxide
Now two hydrogen ions, each lacking an electron and therefore also called protons, attach to the two ends. As we can see in the image below, hydrogen peroxide has no unpaired electrons, so by definition it is not a radical. Nevertheless, as we will explore later, hydrogen peroxide can cause considerable harm in its own way, so it deserves our attention when discussing radicals. Hydrogen peroxide can also be produced through other pathways, but for now, the important point to remember is that it can form from superoxide.
If we represent these steps using Lewis electron-dot structures, the process would look something like the image below.
Singlet oxygen
We have seen that when triplet oxygen, the most common form of oxygen we breathe, gains an electron, it becomes the superoxide anion radical, leaving an unpaired electron in the simplified orbital picture. But imagine a different situation. In triplet oxygen, each oxygen atom is represented as having an unpaired electron in one of its orbitals. Suppose that, under some external influence, one of these unpaired electrons moves into an orbital associated with the other oxygen atom, where it forms a pair, leaving its original orbital empty. This gives us a simplified picture of singlet oxygen. One familiar example of an external influence that can contribute to its formation is ultraviolet light.
Earlier in this article, we saw that during cellular respiration, the energy released as electrons are transferred is used to pump protons from the mitochondrial matrix across the inner membrane. The inner mitochondrial membrane provides an environment in which electrons are transferred one at a time, hydrogen ions are abundant, and oxygen is present. These conditions make it possible for superoxide and hydrogen peroxide to form. The human body is estimated to contain approximately 30 trillion cells.[2] Depending on the type of cell and the tissue in which it is found, a single cell may contain roughly 1,000–2,500 mitochondria.[3] Cellular respiration accounts for approximately 90% of ATP production, and because ATP cannot be stored in substantial quantities for long-term use, it must be produced continuously. So just how much ATP do our mitochondria produce in a single day? An average adult is said to produce and use approximately 100–150 moles of ATP per day, equivalent to roughly 6 × 10²⁵ to 9 × 10²⁵ ATP molecules.[4] Expressed in energy terms, this corresponds to approximately 730–1,090 kcal per day. Is that not an astonishing amount of energy? Once again, I find myself thinking that our bodies are like entire universes in themselves. What matters here is that these figures give us some sense of the enormous scale of ATP production taking place continuously inside our bodies—and, with it, the scale at which electrons may inevitably escape or be misdirected during the process.
So far, we have explored radicals and reactive species associated with oxygen. In the next article, we will take a closer look at how hydrogen peroxide can give rise to the notorious hydroxyl radical.
References
[1] Definition of a radical
[2] Sender R, Fuchs S, Milo R. Revised Estimates for the Number of Human and Bacteria Cells in the Body.
[3] Pizzorno J. Mitochondria—Fundamental to Life and Health.
[4] Physiology, Adenosine Triphosphate
Mechanism of Superoxide and Hydrogen Peroxide Formation by Fumarate Reductase, Succinate Dehydrogenase, and Aspartate Oxidase
