The oxidation number of an element indicates the hypothetical charge of an atom in a compound. It is hypothetical because, in the context of a compound, the elements may not necessarily be ionic. When the number of electrons associated with an atom changes, its oxidation number also changes. When an element loses an electron, its oxidation number increases.
When an element loses an electron, its oxidation number always gets more positive. The exact configuration of oxidation numbers in a compound is specified by a series of oxidation number rules. These rules describe the distribution of oxidation numbers within a compound and outline the typical oxidation numbers for some elements. If you become familiar with these rules, you may be able to understand and predict which reactant will oxidize.
Multiple Oxidation Numbers
Some elements have many possible oxidation numbers. If you know which elements these are, you can predict what will happen to their oxidation numbers in a reaction. For example, iron can have oxidation numbers ranging from -2 to +6. The most common oxidation numbers for iron are +2 and +3. In order to distinguish which of these is present in a compound, scientists write the oxidation state in Roman numerals in the compound name. In the reaction, if the iron loses electrons, its oxidation state will change. This is the case when iron rusts. Solid iron is oxidized to iron (II) by oxygen atoms. Then, the iron (II) atoms lose electrons when reacting with hydrogen ions and oxygen. This reaction forms iron (III) ions, which can go on to form iron (III) hydroxide and iron (III) oxide.
When a compound loses electrons, something must be compelling it to do so. This is called the oxidizing agent. For example, when iron rusts, oxygen is an oxidizing agent. The oxygen receives the electrons that the iron loses. The electrons that are lost in a reaction must be gained somewhere else in order to balance the electric potential. In turn, the oxidation number of the oxygen also changes.
Oxidation and Reduction
Reactions in which an element is oxidized typically involved a corresponding reduction in another element. Reduction happens when an element gains electrons; in this case, its oxidation number is lowered. For example, when iron rusts, oxygen can behave as an oxidizing agent. As the oxygen gains electrons, it changes from an oxidation number of zero to an oxidation number of negative two.
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