About the Limiting Reactant Calculator
This limiting reactant calculator takes a chemical equation and the amounts of the reactants you start with, and tells you which one runs out first. Type the reaction — it does not need to be balanced, because the calculator balances it for you — then enter the mass in grams (or the moles) of each reactant in the order they appear. Molar masses are calculated automatically from the formulas using standard atomic weights.
It is built for chemistry students working stoichiometry problems, lab users planning how much of each reagent to weigh out, and teachers checking answers. Along with the limiting reagent you get the theoretical yield of the product you choose, the extent of reaction, and exactly how much of every excess reactant remains when the reaction stops.
The results assume the reaction goes to completion with no side reactions, which gives the theoretical (maximum) yield. Up to three reactants are supported; element symbols are case-sensitive and ionic charges are not supported.
How to use the limiting reactant calculator
- 1Type the reaction, for example Fe2O3 + CO -> Fe + CO2 — coefficients are optional.
- 2Choose whether you are entering grams or moles.
- 3Enter the amount of each reactant in the order they appear in the equation.
- 4Pick which product you want the theoretical yield for.
- 5Read the limiting reactant, theoretical yield and leftover excess.
Formula and method
First the equation is balanced to get the stoichiometric coefficients aᵢ. Each reactant amount is converted to moles by dividing its mass by its molar mass. Dividing the moles of each reactant by its coefficient tells you how many times the balanced reaction could run on that reactant alone; the smallest of these values, the extent ξ, belongs to the limiting reactant.
The theoretical yield of a product is ξ multiplied by its coefficient (moles), then by its molar mass (grams). Every other reactant is consumed at ξ × its coefficient, and whatever remains is the excess. Comparing masses directly is a common mistake — always compare moles divided by coefficients.
- n
- Moles of a reactant
- M
- Molar mass (g/mol)
- aᵢ
- Coefficient in the balanced equation
- ξ
- Extent of reaction (mol)
Worked examples
Burning hydrogen: 10 g H₂ with 64 g O₂
10 g of H₂ is 4.96 mol (enough for 2.48 runs of 2H₂ + O₂), while 64 g of O₂ is 2.00 mol (2.00 runs). Oxygen runs out first, giving 4.00 mol or 72.06 g of water and leaving about 1.94 g of hydrogen unreacted.
Iron smelting: 100 g Fe₂O₃ with 50 g CO
The balanced equation is Fe₂O₃ + 3CO → 2Fe + 3CO₂. 100 g of Fe₂O₃ is 0.626 mol, but 50 g of CO (1.785 mol) supports only 0.595 runs, so CO is limiting. That makes 1.19 mol (66.46 g) of iron and leaves 4.98 g of Fe₂O₃.
Aluminium chloride from moles: 3 mol Al and 4 mol Cl₂
Balanced, the reaction is 2Al + 3Cl₂ → 2AlCl₃. Aluminium could run 1.5 times but chlorine only 1.333 times, so Cl₂ is limiting. That forms 2.667 mol (355.6 g) of AlCl₃ and leaves 0.333 mol (8.99 g) of aluminium.
Frequently asked questions
What is a limiting reactant?+
The limiting reactant (limiting reagent) is the reactant that is completely used up first. It stops the reaction and therefore sets the maximum amount of product that can form.
How do you find the limiting reactant?+
Convert every reactant to moles, divide each by its coefficient in the balanced equation, and compare. The reactant with the smallest result is limiting.
Is the reactant with the smallest mass always limiting?+
No. Different substances have different molar masses and coefficients. In 2H₂ + O₂ → 2H₂O, 10 g of hydrogen outlasts 64 g of oxygen because hydrogen is so light.
How do I calculate the excess reactant left over?+
Multiply the reaction extent by the excess reactant’s coefficient to find the moles consumed, subtract from the moles you started with, and multiply by its molar mass to convert back to grams.
What is theoretical yield?+
Theoretical yield is the maximum product possible if all the limiting reactant reacts completely. Real (actual) yields are lower; divide actual by theoretical and multiply by 100 to get percent yield.