Oxidation Reduction Reactions
An oxidation-reduction or redox reaction is transformation involving electron transfer. It consists of a half reaction in which a substance loses an electron or electrons (oxidation) and another half reaction in which a substance gains an electron or electrons (reduction).
The substance that gains electrons is the oxidizing agent, while the substance that gives up electrons is the reducing agent. Adding the two half equations algebraically and eliminating the free electrons gives the complete oxidation-reduction equation. This equation, however, is not an assurance that the reaction will proceed spontaneously as written. One way to know the direction of the reaction is to determine the value of the standard free energy change, ΔG0, and then to calculate the free energy change, ΔG, for the complete redox reaction. If the resulting value of ΔG is negative, the reaction will proceed spontaneously as written. If ΔG is positive, the reaction can proceed spontaneously in the opposite direction. If ΔG is zero, the reaction is in equilibrium.
Simple redox reactions occur with the direct transfer of electrons from the reducing agent to the oxidizing agent. An example is the reaction between hydrogen (H2) and chlorine (Cl2) to form hydrogen chloride. In this reaction, hydrogen, the reducing agent, donates two electrons to chlorine which is the oxidizing agent. Because it gains electrons, the oxidizing agent changes in valence and becomes more negative or less positive. In the given example, the valence of chlorine changes from 0 to -1. The reducing agent, on the other hand, loses electrons and hence becomes more positive. Thus hydrogen, being the reducing agent in the given example, changes in valences from 0 to +1.
Some common applications of redox reactions in wastewater treatment are the detoxification of cyanide and the precipitation of chromium. Highly toxic cyanide is converted to nontoxic cyanate and finally to carbon dioxide and nitrogen gas by the action of strong oxidizing agents, such as chlorine gas or sodium hypochlorite. In removing hexavalent chromium, the objective is to reduce it to trivalent chromium, which is less toxic and can be precipitated out in the form of hydroxide. Hexavalent chromium is reduced by sodium bisulfite or ferrous sulfate in an acid medium. Lime can then be applied to precipitate out the trivalent chromium as chromium hydroxide, which collects as a chemical sludge upon settling.
In biochemical redox reactions, however, the electrons go through a series of transfers before reaching the terminal acceptor, which is oxygen in aerobic systems. Examples of these electron carriers are NAD (nicotinamide adenine dinucleotide), NADP (nicotinamide adenine dinucleotide phosphate), and flavoproteins. In anaerobic systems, other materials such as nitrates, sulfates, and carbon dioxide may become the electron acceptor, as in the process of denitrification in which nitrate (NO3) is reductively degraded to molecular nitrogen (N2). In either case, organic or inorganic matter, which serves as food for the microorganisms, is the substance oxidized. A simplified stoichiometry of the complete aerobic biochemical oxidation of glucose, for example, produces six moles of carbon dioxide and six moles of water with the release of 686 kilocalories per mole.
Corrosion and Material Degradation; Electron Acceptor and Donor; Industrial Waste Treatment; Sewage Treatment; Waste Management
Resources
Books
Eckenfelder Jr., W. W. Industrial Water Pollution Control. New York: McGraw-Hill, 1989.
Snoeyink, V. L., and D. Jenkins. Water Chemistry. New York: Wiley, 1980.
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