Redox Reactions and Half-Reactions

Redox reactions and half-reactions

Remember from Chemistry 110 that oxidation and reduction occur simultaneously in a redox reaction. The reactions taking place in electrochemical cells are redox reactions. Two questions should be asked to determine if a reaction is a redox reaction:

• Is there a compound or atom being oxidised?

• Is there a compound or atom being reduced?

If the answer to both of these questions is yes, then the reaction is a redox reaction. For example, the reaction given below is a redox reaction:

\(2\text{Fe}^{3+}(\text{aq}) + \text{Sn}^{2+}(\text{aq})\) \(\to\) \(2\text{Fe}^{2+}(\text{aq}) + \text{Sn}^{4+}(\text{aq})\)

You can write a redox reaction as two half-reactions, one showing the reduction process, and one showing the oxidation process. \(\text{Fe}^{3+}\) is gaining an electron to become \(\text{Fe}^{2+}\). \(\color{red}{\textbf{Iron}}\) is therefore being \(\color{red}{\textbf{reduced}}\) and \(\color{blue}{\textbf{tin}}\) is the \(\color{blue}{\textbf{reducing agent}}\) (causing iron to be reduced). The \(\color{red}{\textbf{reduction half-reaction}}\) is:

\(\color{red}{\textbf{Fe}^{3+}\textbf{(aq) + e}^{-} \to \textbf{Fe}^{2+}\textbf{(aq)}}\)

\(\text{Sn}^{2+}\) is losing two electrons to become \(\text{Sn}^{4+}\). \(\color{blue}{\textbf{Tin}}\) is therefore being \(\color{blue}{\textbf{oxidised}}\) and \(\color{red}{\textbf{iron}}\) is the \(\color{red}{\textbf{oxidising agent}}\) (causing tin to be oxidised). The \(\color{blue}{\textbf{oxidation half-reaction}}\) is:

\(\color{blue}{\textbf{Sn}^{2+}\textbf{(aq)} \to \textbf{Sn}^{4+}\textbf{(aq) + 2e}^{-}}\)

Notice that in the overall reaction the reduction half-reaction is multiplied by two. This is so that the number of electrons gained in the reduction half-reaction match the number of electrons lost in the oxidation half-reaction.