The Role of Pro-Oxidants and Antioxidants in Oxygen Toxicity

Abstract

Both Priestley and Scheele, the independent discoverers of oxygen, recognized the toxic biological effects of oxygen. Scheele (1) actually observed that oxygen inhibited the growth of peas. Priestley (2) remarked that oxygen could burn us up too fast. Soon after, Lavoisier (3) noticed oxygen toxicity in animals, but it was not until almost a century later, in 1878, that the first extensive experiments on oxygen poisoning were reported by Bert (4). In spite of this literature, it is surprising that the deleterious biological effects of oxygen have been generally ignored by biologists. In order to understand the effects of oxygen in living systems, it seems appropriate to consider first some aspects concerning the electronic structure of oxygen. Figure 1 shows the location of the p-electrons for 0, 0-, and 0-. For each of these atomic species, there are, in addition, the s-electrons in the more-stable completed (ls)2 and (2s)2 orbitals. Generally, electrons occupy orbitals as determined by the three following rules: (1) The more-stable orbitals are completed with electrons before the less-stable orbitals. (2) The Pauli exclusion principle requires that two electrons in a filled orbital must have opposite spins resulting in a net spin of zero. (3) The Hund rule requires that two electrons cannot occupy the same orbital until all the orbitals of the same stability have one electron each. In addition, the most-stable states will occur when electrons have the same spin, providing that the Pauli exclusion principle is not violated. Thus, when orbitals of equal stability contain only one electron per orbital, these electrons will have the same spin. It can be seen from Fig. 1 that the 2p orbitals have equal stability, and therefore each of the three 2p orbitals accepts one electron before any one of these orbitals accepts another electron. If an orbital contains only an unpaired electron, then the electron spin results in a magnetic moment, which gives rise to paramagnetism. There is generally a strong tendency to pair electrons in orbitals and eliminate a net electron spin. For this reason free radicals which contain unpaired electrons are usually not stable. Figure 1 illustrates that atomic oxygen and the hydroxyl radical (OH-) contain unpaired electrons and are therefore highly reactive. On the other hand, H20 does not contain unpaired electrons and is relatively stable.