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Redox reaction

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Redox reactions are chemical reactions in which one substance surrenders electrons to another. Redox is the common abbreviated form of the name, which is more formally known as an oxidation-reduction reaction. The reactions that take place in all electrochemical cells, or "batteries," are redox reactions. So, too, are all electrolytic reactions.


Oxidation refers to the loss of electrons. Such a loss always results in an increase in valence or a more positive oxidation state--that is, an adjustment of valence in a positive direction.

Reduction refers to a gain of electrons. Such a gain always results in a decrease (hence "reduction") in valence or a more negative oxidation state--that is, an adjustment of valence in a negative direction.


Reducing agent

A reducing agent, sometimes called a reductant, is a substance that gives electrons to other substances in a redox reaction. Reducing agents, like oxidizing agents, are measured by their strength; a strong reducing agent "forces" the electrons onto the other substance, while a weak reducing agent only gives them away weakly. Reducing agents cause reduction in another substance by giving them electrons; the reducing agent itself is oxidized as it loses electrons.[1]

Oxidizing agent

An oxidizing agent, oxidizer, or oxidant is a substance that can take electrons from other substances. Oxidizing agents can be rated by their strength; strong oxidizing agents take electrons away easily, while weak oxidizing agents take electrons away with difficulty. Most but not all oxidizing agents have oxygen in them; sometimes an oxidizing agent is defined as a substance that gives oxygen to another substance. When an oxidizing agent gains electrons from a reducing agent (a substance that gives up electrons), the oxidizing agent is reduced and the reducing agent is oxidized. Oxidizing agents cause oxidation in other substances by taking their electrons.[1]


The simplest example of an oxidation-reduction reaction is the combination of hydrogen and oxygen to form water. The usual equation describing this reaction is:

2H2 + O2 → 2H2O

Like all gases that are not noble gases, hydrogen and oxygen exist in nature as molecules containing two atoms joined by covalent bonds. Therefore, these two elements must transfer electrons from one to another in order to form a molecule of water. This takes place in two steps.

First, a hydrogen molecule must dissociate into its component atoms--and in the process lose the two electrons that formed the covalent bond between them. Thus we have two hydrogen ions (free protons) and two electrons:

H2 → 2H+ + 2e-

This is the oxidation step (the hydrogen lost the electrons). Notice that the algebraic sum of the valences of all atomic, molecular, or subatomic species must be the same on the reagent side of the equation as on the product side.

Next, in the reduction step, an oxygen molecule must accept four electrons to dissociate into two oxide ions:

O2 + 4e- → 2O2-

To complete the stoichiometric calculation, one must balance the number of electrons released in the oxidation step with the number of electrons taken up in the reduction step. Oxidation and reduction must balance. In this case, one multiplies all the proportions in the reduction step by two, and then adds the two steps together. The numbers of electrons cancel out, and one now has a number of ionic species:

2H2 + 02 → 4H+ + 2O2-

These ions will immediately form covalent bonds, and so one rewrites the product side of the equation thus:

2H2 + O2 → 2H2O

Electromotive Force

All redox reactions are reversible under the appropriate circumstances. However, each reaction will carry with it a certain potential for oxidation or reduction. In general, the stronger the reducing agent (the reagent that becomes oxidized), the weaker the corresponding oxidizing agent that is formed on the other side of an oxidation step. One can measure the potential difference of any given oxidation step and compare it to that of the oxidation step for dissociating hydrogen into hydrogen ions. The potential of a reducing step is the exact algebraic inverse of the potential for an oxidizing step. To compute the electromotive potential for an entire redox reaction, one adds together the potentials for the oxidation step and the reduction step.

Unless that total electromotive potential is positive, the redox reaction will not normally occur. But one can force it to occur, or even force a given redox reaction to occur in reverse, by applying a strong enough electromotive potential, or voltage, between two metal rods or wires inserted either into water or an appropriate solution. Any rod or wire so energized for this purpose is called an electrode (from the Greek elektra amber and hodos a way; from the observation that the ancient Greeks made that amber, when rubbed, gives off sparks). An electrode that tends to produce an excess of electrons is called a cathode (from the Greek kata down, from the lowering of ionic valence), and an electrode that tends to take electrons up is called an anode (from the Greek ana up). (Note that an ion with positive valence is called a cation and one with negative valence is called an anion.)

Even more importantly, one can harness that electromotive force in order to do useful work. The ability to do work is called energy, which comes from the Greek word for "work." This is the basis for the building of electric cells, or "batteries." This also explains why batteries always "wear out"--all that has happened is that the redox reagents are finished reacting.


Redox reactions have a large number of practical applications:

  1. The oxidation and reduction can be separated so the electron flow between the two can be used. This is a battery.
  2. The collection of hydrogen and oxygen from water, either for use in a fuel cell or for another application requiring one or the other of these two gases.
  3. Electroplating, or the deposition of a thin layer of one metal on top of another.


  1. 1.0 1.1 Chemistry:Precision and Design; Verne Biddle and Gregory Parker; A Beka Book

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