More on oxidation/reduction

Electrons move, oxidation states change, and electromotive force can substitute for chemical potential in chemical oxidation and reduction. Of the various principles from physical chemistry that are important, chemical potential is most vital because it underlies the other concepts. Consider the ferrous ion, Fe++. It can be reduced to elemental iron, Fe, or can be oxidized to ferric ion, Fe+++. The reaction will depend on solution chemistry that employs some other ions or molecules to donate electrons or to accept electrons. Oxidation of iron to the ferric state can be very important in pollution control because ferric hydroxide is much less soluble than is ferrous hydroxide. For all reactions, there is a required driving force as an oxidation-reduction potential. Such potentials are relative to other potentials, and different reactions cannot be compared unless there is some reference standard. This standard is by definition the reaction:

H2 = 2 H+ + 2 e¯

This is arbitrarily assigned an electromotive force of zero volts. A very cumbersome and temperamental electrode for this reaction can be fabricated from a catalyst called platinized platinum over which hydrogen in bubbled. The EMF (electromotive force) specification is for hydrogen gas at STP (standard temperature and pressure) that equilibrates with 1 M hydrogen ions. This inconvenient electrode is very seldom used except for laboratory calibrations or with non-aqueous systems where it is necessary to return to the definition of the hydrogen electrode. Hydrogen ion concentration is conveniently expresses as pH. Glass electrodes that are sensitive to hydroxyl ions are used for measuring pH, and the EMF is compared to a reference cell. As anyone who has used a pH meter knows, there are convenient, inexpensive, reliable reference electrodes.

Metals can react with solutions to reduce ions. All solids are assigned an activity of one. Metals tend to displace from solution other metals that have a lower standard potential. However, standard potential is not the whole story. Potential for a given system depends on activity. As with other discussions for environmental engineering, assume that activity and concentration are roughly equal for dilute solutions. However, this assumption is patently false for some of the concentrated solutions encountered in industrial waste treatment. An electrolytic system includes the oxidation reaction and the reduction reaction, usually expressed as half-cells. If the reaction is as written in a table of standard potentials, the sign of the EMF is correct. The sign is reversed if the reaction is reversed. A half-cell is governed by the Nernst equation:


where E = electrochemical potential, volts
Eo = standard potential, volts
R = universal gas constant
T = absolute temperature, ° K
n = number of electrons transferred
F = Faraday's constant, 96,500 coulombs/equivalent
' ox ' and ' red ' are the activities of the oxidized and reduced ionic species.

The concentrations should be molal (moles per liter of water), but there is almost negligible error in using molarity (moles per liter of solution) when the concentrations are small. While oxidation-reduction potential determines the ratio of inorganic ions in solution, it also plays a key role in biochemistry. The terms aerobic and anaerobic are crude statements of oxidation-reduction potential. Some organisms thrive only at high concentrations of oxygen corresponding to high redox potential while many organisms die quickly when even trace concentrations of dissolved oxygen are present and prefer very low oxidation-reduction potential. Metabolism is most energetically efficient when a sequence of steps can make use of enzymatic cofactors that are regenerated at high redox potential. Anaerobic organisms lack these reactions and some have evolved to mechanisms that are poisoned by oxygen.

The important concept for the environmental engineer is that many different biochemical and chemical reactions must interact to develop the oxidation-reduction potential of a solution. The attainment of redox equilibrium is usually rapid, but kinetic effects may be important. Any shift in concentrations for one reaction will affect redox potential and thus will affect all the other concentration ratios. If it were as easy to measure as is pH, redox potential would probably be more common in our arsenal of weapons for monitoring and combatting pollution. Unfortunately, measuring redox potential as the electromotive force of a platinum electrode with a reference electrode has drifting calibration and unreliable readings.

The Nernst equation is also important for many of the special electodes that we use to monitor and measure substances dissolved in water. A common example is the pH meter. The reduced species (water) is essentially constant, and substituting hydrogen ion (the oxidized species) into the equation results in a log term almost the same as the definition of pH.

The straight line tells us that we can use a simple arithmetic scale on a pH meter, but a logarithmic scale would be required if we wished to read hydrogen ion concentration. Please inspect the Nernst equation and decide what the numerical value for the slope of our graph should be. Other electrodes that have slopes that are predicted by this equation are called Nernstian. At extremes of concentration the graphs tend to depart from linearity because the assumption that activity equals concentration is false.

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From H. R. Bungay, BASIC Environmental Engineering, BiLine Associates, 2nd ed , 1992