Oxidation can be defined as the loss of electrons by an atom, ion, or molecule during a chemical reaction. This process is often accompanied by an increase in oxidation state. In terms of electron transfer, oxidation can be represented by the general equation:
\[ \text{Oxidation:} \quad \text{A} \rightarrow \text{A}^{n+} + n \text{e}^- \]
Where:
- \( \text{A} \) is the species being oxidized,
- \( \text{A}^{n+} \) is the oxidized form of the species with a higher oxidation state,
- \( n \text{e}^- \) represents the \( n \) electrons lost.
Example 1: Oxidation of Zinc
A common example of oxidation is the reaction of zinc with copper(II) ions:
\[ \text{Zn} \rightarrow \text{Zn}^{2+} + 2 \text{e}^- \]
In this equation:
- Zinc (Zn) is oxidized to zinc ions (\(\text{Zn}^{2+}\)) by losing two electrons (\(2 \text{e}^-\)).
Example 2: Oxidation of Iron
Another example is the oxidation of iron in the presence of oxygen:
\[ 4 \text{Fe} + 3 \text{O}_2 \rightarrow 4 \text{Fe}^{3+} + 12 \text{e}^- \]
In this reaction, iron (Fe) is oxidized to iron ions (\(\text{Fe}^{3+}\)), losing a total of twelve electrons in the process.
In summary, oxidation involves the loss of electrons and the transformation of a species to a higher oxidation state, as demonstrated in these examples.