Atoms, ions, and molecules that have an unusually large affinity for electrons tend to be good oxidizing agents. Elemental fluorine, for example, is the strongest common oxidizing agent. F 2 is such a good oxidizing agent that metals, quartz, asbestos, and even water burst into flame in its presence. These compounds are strong oxidizing agents because elements become more electronegative as the oxidation states of their atoms increase. Good reducing agents include the active metals, such as sodium, magnesium, aluminum, and zinc, which have relatively small ionization energies and low electro-negativities.
Some compounds can act as either oxidizing agents or reducing agents.
One example is hydrogen gas, which acts as an oxidizing agent when it combines with metals and as a reducing agent when it reacts with nonmetals. Another example is hydrogen peroxide, in which the oxygen atom is in the -1 oxidation state. Because this oxidation state lies between the extremes of the more common 0 and -2 oxidation states of oxygen, H 2 O 2 can act as either an oxidizing agent or a reducing agent. Spontaneous oxidation-reduction reactions convert the stronger of a pair of oxidizing agents and the stronger of a pair of reducing agents into a weaker oxidizing agent and a weaker reducing agent. By convention, all of the half-reactions are written in the direction of reduction.