How to Write Half Equations

A half equation is a chemical equation that shows how one species - either the oxidising agent or the reducing agent - behaves in a redox reaction. If you add two half equations together, you get a redox equation.

Steps

  1. 1
    Identify the species for which you are writing the equation. As an example, let's look at the permanganate ion.
  2. 2
    Start your equation with this species on the left. [1]
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  3. 3
    Identify the species that it is oxidised/reduced to. In our example, permanganate in reduced to manganese ions. Place this species on the right of the equation. [2]
  4. 4
    Identify if it is reduction or oxidation. If it is reduction, you need to place electrons on the left of the equation. If oxidation, on the right. Permanganate is reduced so the electrons go on the left. [3]
  5. 5
    Look at the conditions in which the reaction takes place. Some reductions or oxidations take place under acidic or alkaline conditions. If acidic, you need to add hydrogen ions to the left. If alkaline, add hydroxide ions to the left. In these cases, you usually need to add water to the right as well. In the example, permanganate is reduced under acidic conditions. If you aren't told the conditions it reacts under, here's a rough guide: [4]
    • If the initial species needs to lose oxygen to become the final species (as permanganate does to become manganese), it typically uses acidic conditions.
    • if the initial species needs to gain oxygens to become the final species (such as oxidising vanadium to vanadate), it typically requires alkaline conditions.
    • Otherwise, it probably doesn't require either acidic or alkaline conditions.
  6. 6
    Balance the equation. [5]
    • Balance any metals.
    • Balance any non-metals (except hydrogen and oxygen).
    • Balance oxygen.
    • Balance hydrogen.
    • Use the electrons to balance the charges.
  7. It does get a lot easier with practise, once you've got the hang of it. Try out these ones:
    • Dichromate (Cr2O7[2-]) reduced to chromium ions (Cr[3+]) under acidic conditions.
    • Nitrate (NO3[-]) reduced to nitrogen monoxide (NO) under acidic conditions.
    • Copper (Cu) oxidised to copper ions (Cu[2+])
    • Hydrogen peroxide (H2O2) oxidised to oxygen (O2) under alkaline conditions.
    • Hydrogen peroxide (H2O2) reduced to water (H2O) under acidic conditions.
    • Vanadium ions (V[2+]) oxidised to vanadate (VO4[3-]) under alkaline conditions.
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      References

      1. ↑ https://chem.libretexts.org/Bookshelves/Analytical_Chemistry/Supplemental_Modules_(Analytical_Chemistry)/Electrochemistry/Redox_Chemistry/Balancing_Redox_reactions
      2. ↑ https://www.chemguide.co.uk/inorganic/redox/equations.html
      3. ↑ https://www.bbc.co.uk/bitesize/guides/z9h9v9q/revision/4
      4. ↑ https://chem.libretexts.org/Bookshelves/Analytical_Chemistry/Supplemental_Modules_(Analytical_Chemistry)/Electrochemistry/Redox_Chemistry/Balancing_Redox_reactions
      5. ↑ https://www.chemguide.co.uk/inorganic/redox/equations.html

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