Balancing Redox Equations Using the Half-Reaction Method
Balance the redox reaction.
1. Assign oxidation states to all atoms and identify the substances being oxidized and reduced.
2. Separate the overall reaction into two half-reactions, one for oxidation and one for reduction.
OXIDATION \mathrm{Al}(s)\longrightarrow\mathrm{Al}^{3+}(a q)
REDUCTION {\mathrm{C u}}^{2{+}}(a q)\longrightarrow\mathrm{Cu}(s)
3. Balance each half-reaction with respect to mass in the following order:
• Balance all elements other than H and O.
• Balance O by adding \mathrm{H}_{2}\mathrm{O}.
• Balance H by adding \mathrm{H}^{+}.
All elements other than hydrogen and oxygen are balanced, so you can proceed to the next step.
No oxygen; proceed to the next step.
No hydrogen; proceed to the next step.
4. Balance each half-reaction with respect to charge by adding electrons to the right side of the oxidation half-reaction and the left side of the reduction half-reaction. (The sum of the charges on both sides of each equation should be equal.)
\mathrm{Al}(s)\longrightarrow\mathrm{Al}^{3+}(a q) +3\mathrm{e^{-}}2\mathrm{e}^{-}+\mathrm{Cu}^{2+}(a q)\,\longrightarrow\,\mathrm{Cu}(s)
5. Make the number of electrons in both half-reactions equal by multiplying one or both halfreactions by a small whole number.
2\ \times\mathrm{[Al(s)}\longrightarrow\mathrm{Al}^{3+}(a q)\ +3\mathrm{e}^{-}]3\,\times\,[2e^{-}\,+\,\mathrm{Cu}^{2\,+}(a q)\,\longrightarrow\,\mathrm{Cu}(s)]
6. Add the two half-reactions together, canceling electrons and other species as necessary.
2\,\mathrm{Al}(s)\longrightarrow2\,\mathrm{Al}^{3+}(a q)+\cancel{6\mathrm{e}^{-}}{\underline{{\cancel{6~\mathrm{e}^{-}}}+3\,{\mathrm{Cu}}^{2+}(a q)\longrightarrow3\,{\mathrm{Cu}}(s)~~~~~}}
2\,\mathrm{Al}(s)+3\,\mathrm{Cu^{2+}}(a q)\,\longrightarrow
2\,\mathrm{Al}^{3+}(a q)+3\,\mathrm{Cu}(s)
7. Verify that the reaction is balanced with respect to both mass and charge.
Reactants Products
2 Al 2 Al
3 Cu 3 Cu
+6 charge +6 charge