Oxygen. It’s a matter of life and death. Without oxygen most life on earth would cease to exist. But the same reactivity that makes it ideal for using sugars for energy can also backfire and damage the very cells that are using that energy. If allowed to run amok, oxygen can damage DNA, lipids, or proteins. Damage to DNA will either kill cells or turn cancerous. Damage to lipids causes damage to membranes. Damage to proteins inactivates enzymes. All of these effects result in serious disease.
Over time, evolution has developed ways to help mitigate this problem. Antioxidants are compounds that help protect cells from oxidative damage. They function by blocking the reactive compounds caused by oxygen. A quick glance at a drug store’s health food area will demonstrate the tremendous growth in popularity of antioxidants.
Although the popular press provides a lot of recommendations and superficial second-hand information about antioxidants, many of these reports misinterpret original research findings. It is only by referring to primary research journals critically that we can find the truth about antioxidants.
One particular class of antioxidants is the group of phenolic antioxidants. They are so called because they are based on phenol, an alcohol composed of a benzene ring and a hydroxyl group. These are particularly interesting because as recently as a few years ago, they were believed only to be important for flavour (Escarpa and Gonzales 2001).
Any high school chemistry student knows that molecular oxygen comprises two oxygen atoms joined by a double-bond. Empirical evidence, however, begs to differ. The length of the bond between the oxygen atoms has been shown to be shorter (and stronger) than a single bond, yet longer (and weaker) than a double bond (Levine and Kidd, 1985; Halliwell, 1984). In 1980, Demopoulos et al. (1980) showed that molecular oxygen is in fact a diradical—that is, it has a resonance structure with one single bond and two unpaired electrons in different orbitals.
Forman and Fisher (1980) further refined the model to illustrate that ground state oxygen contains two unpaired electrons with same spins in separate orbitals, and is thus paramagnetic:
Figure 1 Ground state oxygen. Source: Levine and Kidd, 1985. p. 21
Consequently, oxygen is much more reactive than would be expected by its high school definition.
Given its tremendous reactivity, that life on earth evolved to use it is clearly a Catch-22—most life requires oxygen to live; yet, the same oxygen can easily run amok and oxidize vital cellular structures. This fact has given rise to the free radical theory of aging, which states that almost all aging processes are directly related to damage by free radicals (Atkins, 2000; Perez-Campo, 1998). In most of these cases, the myriad species of free radicals can all be traced back to chain reactions initiated by reactive oxygen species.
While all forms of oxygen are very reactive—hence the term “oxidation” for all losses of electrons—there are four forms of oxygen that are particularly reactive. These reactive oxygen species include—in order of increasing reduction—singlet oxygen, superoxide radical, hydrogen peroxide, and hydroxyl radical.
Once biomolecules are damaged, they can lead to aging, cancer, or even heart disease and stroke.
So-called because of its single peak on an Electron Spin Resonance graph (Halliwell, 1984), singlet oxygen can exist in two forms. One form is an energetic form of molecular oxygen wherein the two outer shell electrons share the same orbital. This form is therefore diamagnetic and not a radical. The other form has the two outer shell electrons in different orbitals, as in ground state oxygen; however, the electrons are of opposite spins. This form of singlet oxygen is therefore quasi-radical (Halliwell, 1984). It is capable of rapidly oxidizing many molecules, including membrane lipids. Halliwell and Gutteridge (1984) point out that single oxygen is commonly produced in human eye lenses. This species could therefore be important in the formation of cataracts and age-related hardening of the lens (presbyopia).
Figure 2 Singlet oxygen electron configuration. Source: Levine and Kidd, 1985
O2 + e- •O2-
Adding a single electron to molecular oxygen produces superoxide radical (Halliwell, 1984), which is an anion due to its net negative charge. Although in aqueous solutions, it is only moderately reactive, it is extremely reactive in organic solvents. It is largely responsible for causing ordinarily-unreactive carbon tetrachloride (CCl4) to become toxic by forming chlorine radicals. In aqueous solutions, this radical slowly combines with another superoxide radical to form one molecule of hydrogen peroxide (H2O2) (Halliwell, 1984).
Figure 3 Superoxide electron configuration. Source: Levine and Kidd, 1985
Superoxide radicals are cytotoxic, fragmenting DNA, causing inflammation, and damaging cellular structures. Halliwell and Gutteridge (1984) also point out that helpful cells, such as immune system cells, produce superoxide radical to bombard and destroy foreign invaders.
