How to Balance a Chemical Equation
Balancing means making the number of atoms of each element the same on both sides, because atoms are neither created nor destroyed in a reaction. You change only the coefficients in front of formulas — never the subscripts inside them, because that would turn the substance into something else entirely.
Why we balance at all
A chemical equation is a statement of the law of conservation of mass: the same atoms that go in must come out, only rearranged. An unbalanced equation describes a reaction that breaks that law, so it cannot be used for any calculation. Once balanced, the coefficients double as the mole ratios that drive every stoichiometry problem.
The method
- Write the unbalanced equation and list how many atoms of each element sit on each side.
- Balance one element at a time, starting with an element that appears in only one compound on each side.
- Leave oxygen and hydrogen for last. They usually appear in several compounds, so they fall into place once everything else is set.
- Adjust coefficients only. If a fraction appears, multiply the whole equation through to clear it.
- Check every element, and check that total charge balances too if ions are involved.
Worked example: combustion of propane
C₃H₈ + O₂ → CO₂ + H₂O.
- Carbon: 3 on the left, so put 3 in front of CO₂.
- Hydrogen: 8 on the left, so put 4 in front of H₂O (4 × 2 = 8).
- Oxygen last: the right now has 3×2 + 4×1 = 10 O atoms, so put 5 in front of O₂.
Every element matches, so the equation is balanced. Combustion problems almost always work cleanly when carbon goes first, hydrogen second and oxygen last.
The polyatomic-ion shortcut
When a polyatomic ion such as sulfate (SO₄²⁻) or nitrate (NO₃⁻) stays intact on both sides, balance it as a single unit instead of counting S and O separately.
For Al + H₂SO₄ → Al₂(SO₄)₃ + H₂, treat SO₄ as one block:
Three sulfate blocks on each side, two aluminium, six hydrogen — balanced in far fewer steps than tracking nine oxygen atoms by hand.
When a fraction appears
Half-coefficients are common in combustion. For C₂H₆ + O₂ → CO₂ + H₂O, balancing C and H first leaves O₂ at 3½:
×2 → 2 C₂H₆ + 7 O₂ → 4 CO₂ + 6 H₂O
Multiplying every coefficient by 2 clears the fraction and gives the conventional whole-number form.
Common mistakes
- Changing a subscript to balance — never do this, as it makes a different compound.
- Forgetting a coefficient multiplies every atom in the formula, not just the first one.
- Leaving a fraction in the final answer instead of scaling the whole equation up.
- Ignoring charge in ionic equations — both mass and charge must balance.
Once the equation is balanced, the coefficients are your mole ratios. Carry the known amount through with the Stoichiometry Calculator, handle two reactants with the Limiting Reagent Calculator, and see how it all fits together in How to Approach Stoichiometry.
The General Chemistry Workbook works through balancing, mole ratios and yields with a full answer key.
Frequently Asked Questions
No. Subscripts define the substance, so changing one creates a different compound. You may only place coefficients in front of whole formulas. Changing H₂O to H₂O₂ would turn water into hydrogen peroxide.
Oxygen and hydrogen usually appear in several compounds at once, so their counts keep shifting while you adjust other elements. Saving them for last means they fall into place after everything else is fixed.
Finish balancing with the fraction in place, then multiply every coefficient in the equation by the smallest number that clears it, usually 2. This keeps the ratios correct while giving a whole-number answer.
If a polyatomic ion such as sulfate or nitrate stays intact on both sides, balance it as a single unit rather than counting its atoms separately. This is faster and less error-prone.
Because mass is conserved: the atoms that react must all appear in the products. A balanced equation also gives the mole ratios used in every stoichiometry calculation, so an unbalanced one cannot be used.