The concept of oxidation arises from the combination of elemental oxygen with other elements to form oxides, as in this example using aluminum:
4Al + 3O 2 → 2Al 2O 3
With oxidation numbers inserted as superscripts, this reaction is written
to show that both elements change oxidation numbers. Because the oxidation numbers changed, an oxidation‐reduction reaction is defined as one in which electrons are transferred between atoms. In the example, each oxygen atom has gained two electrons, and each aluminum has lost three electrons.
In an electron transfer reaction, an element undergoing oxidation loses electrons, whereas an element gaining electrons undergoes reduction. In the aluminum‐oxygen example, the aluminum was oxidized, and the oxygen was reduced because every electron transfer reaction involves simultaneous oxidation and reduction. These reactions are frequently called redox reactions.
A subtlety deserving your close attention is that the oxidizing agent (in the example, oxygen) is reduced, whereas the reducing agent (in the example, aluminum) is oxidized.
Because chemists have defined oxidation in terms of electron transfer, it is quite unnecessary for redox reactions to have oxygen as the oxidizing agent. Study the next example of metallic zinc reacting with chlorine gas to form zinc chloride:
The oxidizing agent that gains electrons is chlorine, and the reducing agent that loses electrons is zinc.
A valuable generalization is that the nonmetals in the upper right region of the periodic table are strong oxidizing agents. The metals in their elemental state are strong reducing agents, as is hydrogen gas.
- In the following redox reaction, identify the element that is oxidized, the element that is reduced, the oxidizing agent, and the reducing agent.
