It is important to understand the properties of dioxygen since oxidation reactions using it power not only our bodies but our entire civilization. We will obviously concentrate on biological reactions, but even these show the same characteristics as non-biological ones.
- oxidation of organic molecules by oxygen is thermodynamically favored but kinetically slow.
- pure oxygen environments are toxic to cells and organisms.
First we will try to understand these properties of oxygen, and then we will see how organisms overcome these problem to use dioxygen.
We can understand both of these properties by looking at the molecular orbitals of oxygen and its reduction products as shown in the diagrams below. Ground state oxygen is a diradical, which explains the paramagnetic behavior of oxygen. The two unpaired oxygens each have a spin state of 1/2 for a total resultant spin S of 1, making ground state oxygen a triplet (2S+1) = 3. Organic molecules typically undergo 2 electron oxidation steps. Consider the stepwise oxidation of methane below. The oxidation number of C in methane is -4, -2 in methanol, 0 in formaldehyde, +2 in formic acid, and finally +4 in carbon dioxide, indicating two electron losses in each step.
The two electrons lost by the organic substrate are added to oxygen, but since the two lost electrons are spin paired, a spin flip must occur to allow the electrons to enter the unfilled oxygen orbitals. Alternatively, energy can be put into ground state dioxygen to produce excited state singlet oxygen (S=0, 2S+1 = 1). The source of the large activation energy required (about 25 kcal or 105 kJ/mol) to flip the electron spin accounts for the kinetic sluggishness of reactions of dioxygen with organic reactants.
A traditional Lewis structure for ground state dioxygen can not be easily written since the electrons are added in pairs, and dioxygen is a diradical. There are 6 electrons in the sigma molecular orbitals from second shell electrons (two each in σ2s, σ2s*, and σ2p,) and 6 electrons in the pi molecular orbitals from second shell electrons (two each in two different π2p orbitals, and one electron each in two different π2p*), so the net number of electrons in bonding orbitals is 4, giving a bond order (or number of 2). In contrast it is easy to write the Lewis structure of singlet, excited state oxygen, since all electrons can be viewed as paired, with two net bonds (1 sigma, 1 pi) connecting the atoms of oxygen. This Lewis structure will be used to represent singlet, excited oxygen, which should react more quickly with organic molecules. The excited state single on the far right (below) is unstable and decays to the middle singlet state. The middle state is approximately 94.3 kJ/mol higher in energy than the ground state triplet (on the left). In quantum mechanical parlance, the transition from the ground state triplet to the singlet state is forbidden for a number of reasons, making it unlikely that absorption of a photon will induce the transition.
The Reductions of Dioxygen
When oxygen oxidizes organic molecules, it itself is reduced. By adding electrons one at a time to the molecular orbitals of ground state dioxygen we produce the step-wise reduction products of oxygen. On the addition of one electron, superoxide is formed. A second electron produces peroxide. Two more produces 2 separated oxides since no bonds connect the atoms (the number of electrons in antibonding and bonding orbitals are identical). Each of these species can react with protons to produce species such as HO2, H2O2 (hydrogen peroxide) and H2O. It is the first two reactive reduction products of dioxygen that make it potentially toxic.
How are the potential problems in oxygen chemistry dealt with biologically?
Kinetic sluggishness: Enzymes that utilize dioxgen must activate it in some way, which decreases the activation energy. Enzymes that use dioxygen typically are metalloenzymes, and often heme-containing proteins. Since metals such as Fe2+ and Cu2+ are themselves free radicals (i.e. they have unpaired electons), they react readily with ground state oxygen which itself is a radical. The molecular orbitals of the metal and oxygen combine to produce new orbitals which for oxygen are more singlet-like in nature. Likewise, dioxygen reacts more readily with organic molecules which can themselves form reasonably stable free radicals, such as flavin adenine dinucleotide (FAD), as we shall see later.
Dioxygen toxicity: Since toxicity arises from the reduction products of oxygen, enzymes that use oxygen have evolved to bind oxygen and its reduction products tightly (through metal-oxygen bonds) so they are not released into the cells where they can cause damage. In addition, enzymes which detoxify free dioxygen reduction products are widely found in nature. For example:
- superoxide dismutase catalyzes the dismutation (self-redox) of 2 superoxides into dioxygen and hydrogen peroxide;
- catalase converts hydrogen peroxide into water and oxygen;
- peroxidase catalyzes the reaction of hydrogen peroxide with an alcohol to form water and an aldehyde
- peroxiredoxins react with peroxides and thioredoxin (a small electron donor) to form water and oxidized thioredoxin .
Finally free radical scavengers such as vitamins A, C, E, and selenium can react with reactive free radicals to produce more stable free radical derivatives of the vitamins and Se. More on this later.