Redox Half-Reaction Balancer (Ion-Electron Method)
Enter an unbalanced redox half-reaction like MnO4- -> Mn2+ and this tool automatically balances it using the ion-electron (half-reaction) method, adding H+, H2O, OH-, and electrons (e-) as needed. Supports both acidic and basic conditions.
Examples of common redox half-reactions
Half-reactions that frequently appear in high school and college chemistry, along with their balanced results. Use them to check your own calculations.
| Half-reaction | Condition | Unbalanced half-reaction | Balanced half-reaction |
|---|---|---|---|
| Reduction of permanganate ion (to Mn2+) | Acidic (balance with H+) | MnO4- → Mn2+ | MnO4- + 8H+ + 5e- → Mn2+ + 4H2O |
| Reduction of dichromate ion (to Cr3+) | Acidic (balance with H+) | Cr2O7^2- → Cr3+ | Cr2O7^2- + 14H+ + 6e- → 2Cr3+ + 7H2O |
| Oxidation of iron(II) ion (to iron(III) ion) | Acidic (balance with H+) | Fe2+ → Fe3+ | Fe2+ → Fe3+ + e- |
| Reduction of chlorine (to chloride ion) | Acidic (balance with H+) | Cl2 → Cl- | Cl2 + 2e- → 2Cl- |
| Oxidation of sulfite ion (to sulfate ion) | Acidic (balance with H+) | SO3^2- → SO4^2- | SO3^2- + H2O → SO4^2- + 2H+ + 2e- |
| Reduction of nitrate ion (to nitric oxide) | Acidic (balance with H+) | NO3- → NO | NO3- + 4H+ + 3e- → NO + 2H2O |
| Reduction of permanganate ion (to manganese dioxide, basic) | Basic (balance with OH-) | MnO4- → MnO2 | MnO4- + 2H2O + 3e- → MnO2 + 4OH- |
Tips
- Ion charges are recognized just by writing them at the end, like "Fe2+" or "Cl-". For polyatomic ions such as SO4^2-, use caret notation (^2-) to avoid misparsing.
- Choosing the basic condition shows the result after first solving under acidic conditions and then converting with OH-, so you can follow the same procedure as in a textbook.
- Reading the "Balancing steps" panel lets you trace the half-reaction method exactly: balance the central element, then oxygen, then hydrogen, then charge.
- This tool only supports half-reactions with a single central element (any element other than H or O). Complex equations where multiple elements change oxidation state are outside its scope.
- Load a sample half-reaction from the buttons first to see the result layout, then try entering your own half-reaction.
Frequently Asked Questions
Side Note — Half-reactions and how batteries work
The concept of a half-reaction is more than just a coefficient-balancing trick — it's also the foundation for understanding how batteries (voltaic cells, Daniell cells, dry cells, and more) work. Inside a battery, an oxidation half-reaction (releasing electrons) proceeds independently at the negative electrode while a reduction half-reaction (accepting electrons) proceeds at the positive electrode; current flows as those electrons travel through the external circuit. In other words, the half-reactions this tool calculates are exactly what's happening at one electrode of a battery.
Half-reactions involving the permanganate ion (MnO4-) are a staple of redox titrations in chemistry labs because the color change is so easy to see. The deep purple MnO4- being reduced to pale pink (or nearly colorless) Mn2+ has long served as its own built-in indicator for spotting the titration endpoint, without needing a separate indicator dye.
The reason hydroxide ions (OH-) show up in half-reactions under basic conditions is that, in an environment where the concentration of H+ is extremely low (effectively negligible), writing H+ directly into the equation would be chemically unnatural. Textbooks usually explain this conversion as "solve it as acidic first, then rewrite with OH-," and this tool's calculation procedure is essentially that same textbook method turned into an algorithm.