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.

Last updated: 2026-09-20

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-

Balancing half-reactions by the ion-electron method

Balancing a redox equation by inspection tends to run into a wall. **The ion-electron method** divides the reaction into an oxidation half and a reduction half and balances each in a fixed order. The procedure has four steps: **balance the central element, balance oxygen with water, balance hydrogen with protons, then balance charge with electrons.**

**Basic conditions add a further step.** After balancing exactly as under acidic conditions, add as many hydroxide ions to both sides as there are protons, and combine each resulting proton and hydroxide into water. Omitting this leaves protons sitting in a basic solution, which is not a sensible equation. Enter an unbalanced half-reaction such as `MnO4- -> Mn2+` and this tool supplies the protons, water, hydroxide and electrons the chosen conditions require and returns the balanced equation.

How to balance a half-reaction

Enter the half-reaction The arrow may be written `->`, `=` or `→`.
Write ionic charges correctly Place them at the end, as in `Fe2+`, `Cl-` or `SO4^2-`. Caret notation is understood.
Choose the conditions **Acidic conditions balance with protons, basic conditions with hydroxide.**
Read the result The balanced equation appears with the numbers of water, protons, hydroxide and electrons added.

Tips for getting more out of it

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.

Where this helps

Checking chemistry homework

Compare an equation you balanced by hand against the result.

Understanding acidic against basic

Running the same half-reaction under both conditions makes the purpose of the hydroxide step clear.

Assembling the overall equation

**Equalise the electrons in the two halves and add them, and you have the overall equation.** This is the step before that.

Thinking about cells and electrolysis

What happens at each electrode is itself a half-reaction, so you can work through them electrode by electrode.

Redox terms explained

Half-reaction: An equation for the oxidation or the reduction alone. **Electrons appearing in the equation is what distinguishes it from the overall one.**
Oxidation: The loss of electrons. In a half-reaction **the electrons appear on the right.**
Reduction: The gain of electrons. In a half-reaction the electrons appear on the left.
Ion-electron method: Balancing by splitting into half-reactions and proceeding in order. **It is more mechanical than the oxidation-number method.**
Acidic conditions: Balancing with protons and water.
Basic conditions: **Balance as under acidic conditions, then add hydroxide to both sides to cancel the protons.**

Frequently Asked Questions

Under acidic conditions, any excess or shortage of oxygen is balanced with water (H2O), and hydrogen with hydrogen ions (H+). Under basic conditions, the same H+ amount is first worked out as if the reaction were acidic, then that many hydroxide ions (OH-) are added to both sides; on the side that had H+, the H+ and OH- combine into H2O, and finally any water common to both sides is cancelled out to reach the final form.

In a redox reaction, electrons are transferred between the species being oxidized (which loses electrons) and the species being reduced (which gains them). Since a half-reaction represents just one half of that overall reaction, the electron transfer must be written explicitly as e- to make the charge balance work out.

No. This tool is scoped to the common textbook pattern where only one element (the central element) changes oxidation state. Complex equations, such as redox reactions involving organic compounds where multiple elements change oxidation state at once, may not be solved correctly.

MnO4- contains 4 oxygen atoms, so 4 water molecules must be added to the product side to account for them. Those 4 water molecules contain a total of 8 hydrogen atoms, so 8 H+ ions must be added to the reactant side to match that hydrogen count.

The half-reaction method splits a complex redox equation into two separate half-reactions — oxidation and reduction — balances each independently, and then combines them once the electron counts match. It's easier to follow than balancing the whole equation at once, and it especially helps organize how H+, OH-, and H2O should be handled under acidic versus basic conditions.

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.

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