Figure 6.2 Movement of electrons in an oxidation-reduction reaction.
The movement of electrons can be tracked by splitting this reaction into two half-reactions. The electrons, symbolized as e-, have to move from the chemical bond holding the two hydrogen atoms together (oxidizing the hydrogen):
to the oxygen to form a new chemical bond (reducing the oxygen):
Where the electrons will go can be judged by comparing how strongly each bond, the hydrogen-hydrogen bond of H2 or the oxygen-hydrogen bond of H2O, attracts the electrons. This is quantified by the redox potential, symbolized as E'o, of each of the half-reactions:
The oxygen-hydrogen bond has a stronger affinity for electrons than the hydrogen-hydrogen bond.2 To esti-
2. The negative redox potential for hydrogen-hydrogen bonds does not mean that electrons are repelled from hydrogen-
mate the potential energy driving the electrons from one bond to the other, calculate the difference in redox potential between the electron acceptor (in this case, the hydrogen-oxygen bond) and the electron donor (in this case, the hydrogen-hydrogen bond):
So, there is a potential energy difference (1.24 volts, to be precise) that will drive electrons from H—H bonds to O—H bonds. The reaction will therefore proceed spontaneously: all that is needed is to bring the two molecules together closely enough and keep them there long enough for the electrons to make the move. The reverse reaction, namely getting hydrogen and oxygen away from water (essentially moving electrons from oxygen-hydrogen bonds to hydrogen-hydrogen bonds) will not be spontaneous. Here, the donor and acceptor bonds are switched around, and the potential energy difference is now -0.42 — 0.92 = —1.24 V. If the reaction is to proceed, work must be done to make it go. In photosynthesis, obviously, the energy to do this work comes from the capture of light.
The question remains, though: what is so bad about oxygen? Ironically, the very thing that makes it so toxic is what makes it so useful to organisms like us. To see why, let us look into the role oxygen plays in the metabolism of glucose. I have already outlined the basic reaction in Chapter 2. To refresh your memory, glucose is combined with oxygen to produce water, carbon dioxide, and energy:
C6H12O6 + 6O2 ^ 6CO2 + 6H2O + 2.82 MJ [mol glucose]-1
hydrogen bonds. Redox potentials are referenced to a standard voltage, in the same way voltages in electronic devices are referenced to a ground. In the case of the oxidation of hydrogen, the negative value simply means the potential has a voltage lower than the agreed-upon reference value.
This simple equation masks a rather complicated set of chemical reactions (Fig. 6.3) that dismantles glucose piece by piece, if you will, and channels the electrons through mediators that capture the energy in them and stores it in ATP. Oxygen's part in all this is to serve as an ultimate acceptor of electrons released from the carbon-hydrogen bonds in glucose, incorporating them into the oxygen-hydrogen bonds of water. Along the way, the electrons released do chemical work for the cell. Because electrons are drawn so powerfully to oxygen, the work they can do is prodigious. The advantage is made clear by comparing how much energy can be captured from glucose when oxygen is present and when it is not. In anaerobic conditions (oxygen absent), only about 7 percent of the energy in glucose is converted to ATP. The remainder is left behind in the bonds of the lactate. When oxygen is present—that is, under aerobic conditions—the energy yield goes up nearly seven-fold, to roughly 40 percent (Fig. 6.3). Under the right conditions, it can exceed 95 percent. That kind of yield is not small potatoes.
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