Limiting Reactant Explained for Stoichiometry

In a chemical reaction, reactants combine in fixed mole ratios set by the balanced equation. If you start with unequal mole amounts relative to those ratios, one reactant will be consumed completely before the others — that is the limiting reactant. It determines the maximum amount of product you can form.

Step-by-step identification method

1. Write the balanced chemical equation
2. Convert each reactant mass (or volume, for gases) to moles
3. Divide each mole amount by its coefficient in the equation
4. The smallest ratio identifies the limiting reactant
5. Use the limiting reactant's moles to calculate product moles via equation ratios

Worked example

Consider 2 H₂ + O₂ → 2 H₂O with 4.0 g H₂ (2.0 mol) and 20.0 g O₂ (0.625 mol). Mole ratios divided by coefficients: H₂ → 2.0/2 = 1.0; O₂ → 0.625/1 = 0.625. Oxygen is limiting. Maximum water: 0.625 mol O₂ × (2 mol H₂O / 1 mol O₂) = 1.25 mol H₂O, or about 22.5 g.

Excess reactant

Reactants not fully consumed are in excess. After the reaction stops, excess reagent remains in the product mixture. Calculating how much excess remains requires subtracting the amount that reacted (determined by the limiting reactant's mole ratio) from the starting amount.

Link to percent yield

Theoretical yield is the maximum product predicted from the limiting reactant. Actual yield from an experiment is often lower due to side reactions, incomplete conversion, or loss during purification. Percent yield = (actual / theoretical) × 100%. The Percent Yield Calculator compares your lab result to the stoichiometric maximum.

Common pitfalls

• Comparing gram amounts directly instead of moles
• Using an unbalanced equation for mole ratios
• Identifying the reactant with the smaller mass as limiting — mass alone does not determine this
• Forgetting to convert all reactants to moles before comparing

Frequently asked questions