Oxidation & Reduction

In the preceeding section we saw that half reactions are reported in tables with E°' values given in Volts. E°' is known as the reduction potential, which is a measure of the tendency of a compound to gain electrons. The reduction potentials of elements and compounds are determined experimentally.

Half reactions: Reduction or oxidation?

Reduction potentials are useful because they indicate how likely a half reaction will take on the role of the “reduction” in a redox reaction. When two half reactions are paired in a redox reaction, the half reaction with the higher E°' will act as the reduction reaction, and the half reaction with the lower E°' will act as the oxidation reaction. Consider the words “reduction potential”—the higher the E°' value, the stronger the tendency for the compound to gain electrons.

In other words, electrons flow from the substance whose half reaction has the lower reduction potential to the substance whose half reaction has the higher reduction potential.

The difference in the E°' values of different half reactions has very real consequences for life on our planet. For example, the half reaction for the reduction of oxygen to water has a very high E°' value, 0.816 V.

It is no coincidence that oxygen has a very high E°' and that we require it for life. Oxygen’s great affinity for electrons allows it to be used in living organisms as an electron “sink” or dumping ground for excess electrons generated by biochemical reactions needed for life. Without the presence of oxygen in our cells to sweep away these excess electrons, aerobic life would soon cease.

To better understand the utility of E°', let’s revisit a reaction we saw earlier:

We know that the two corresponding half reactions from the half reaction table are

Both are written as reduction reactions, yet we know that when coupled together in a redox reaction, one of these reactions will be driven backwards to act as the oxidation reaction. In section 2 of this review, we reversed reaction (1) and added (reverse 1) to reaction (2) to arrive at the overall redox reaction as written above. However, remember that a chemical reaction such as the one above can proceed in either direction.

Which of the half reactions above is more likely to act as the reduction half reaction, and which is more likely to act as the oxidation half reaction? In which direction is the overall reaction most likely to proceed? To answer those questions, remember the rule that a higher E°' value indicates a stronger tendency for the compound to gain electrons. Half reaction (1) has a higher E°' value than half reaction (2), and thus is more likely to act as the reduction (remember OIL RIG):

Adding the two half reactions yields the overall redox reaction, written so that the reaction will spontaneously occur from left to right:

The following value reflects the tendency of a chemical reaction to proceed in the direction written, called the DE°' of the overall reaction:

First look again at the two half reactions, written as reductions, from the half reaction table:

Because reaction (1) has the higher E°' value, it will act as the reduction reaction. Reaction (2) will therefore act as the oxidation reaction. Then, plug the E°' values into the equation:

The tendency for a reaction to proceed in the direction it is written can be determined from DE°': A positive DE°' indicates that the reaction will proceed in the direction it is written. So for our reaction, it will proceed from left to right as below:

You might note that this reaction is the reverse of the reaction we started with, which is the reaction catalyzed by succinate dehydrogenase of the citric acid cycle. If the reverse of the citric acid cycle reaction has the higher DE°', then this citric acid cycle reaction is not favored (not spontanous), and requires energy input. We will look at free energy and spontaneous redox reactions in the next section.

Example 4: Another Approach to Calculating the Change in Reduction Potential (DE°')

The DE°' can also be found by adding the two half reactions in the same way that algebraic equations can be added. Finding the DE°' is then as simple as adding the E°' values of the half reactions.

Look again at the reaction for the conversion of pyruvate into lactate:

pyruvate + NADH + H⁺ lactate + NAD⁺

The two reduction half reactions are as follows

(1)	pyruvate + 2 H⁺ + 2 e^– lactate	E°' = –0.190 V
(2)	NAD⁺ + 2H⁺ + 2e^– NADH + H⁺	E°' = –0.320 V

In order to get the largest value for the DE°' (which is the spontaneous direction), reaction 2 must be written in reverse. When a reduction reaction is reversed to create an oxidation half reaction, the sign of the reduction potential must also be reversed (to get an oxidation potential):

(2 reversed)

NADH + H⁺

NAD⁺ + 2H⁺ + 2e^–

E°' = 0.320 V

Then, reaction 1 and the reversed reaction 2 can be added together to arrive at the net redox reaction, and the DE°' can be found by adding the reduction potential of the first half-reaction to the oxidation potential of the second half reaction. Note that when we look at reactions this way, we do not need to remember which half reaction one gets subtracted from the other (for it doesn't matter when we add).

(1)	pyruvate + 2 H⁺ + 2 e^– lactate	E°' = –0.190 V
(2 reversed)	NADH + H⁺ NAD⁺ + 2H⁺ + 2e^–	E°' = 0.320 V

(sum)	pyruvate + NADH + H⁺ lactate + NAD⁺	DE°' = 0.130 V

Since the DE°' is greater than 0 and is larger than if we did it the other way and reversed reaction 1, the reaction will proceed as written.

[Note–if we instead reversed half reaction 1 and left half reaction 2 as is we would get: 0.190 + (–0.320) = –0.13. This is less than above.]

Redox Reactions

REDUCTION POTENTIAL

Half reactions: Reduction or oxidation?

Example 4: Another Approach to Calculating the Change in Reduction Potential (DE°')