how-to · 5 min read

Limiting reactants and percent yield, with worked examples

Two reactants mixed in arbitrary amounts rarely match the balanced ratio. One runs out first, the reaction stops, and the other is left over. Finding which one limits the reaction is the first step in predicting how much product can form.

This guide works through the limiting reactant comparison, the theoretical yield calculation and the percent yield that compares prediction with what a lab actually collects. The numbers follow the same steps as the limiting reactant and theoretical yield calculators on this site.

By FormatFlow EditorialUpdated October 21, 2026More Chemistry Tools →

Finding the limiting reactant

Convert every reactant amount to moles first. For the reaction 2H2 + O2 → 2H2O, 4 g of H2 is 4 / 2.016 = 1.984 mol, and 32 g of O2 is 32 / 31.998 = 1.0001 mol. Comparing masses directly would be misleading because hydrogen is much lighter per mole, so the mole conversion is essential.

Divide each amount by its coefficient. Hydrogen gives 1.984 / 2 = 0.9921, and oxygen gives 1.0001 / 1 = 1.0001. The smaller quotient is hydrogen, so H2 is the limiting reactant and oxygen is present in excess. The reaction consumes hydrogen twice as fast as oxygen, which is what the coefficient 2 represents.

The same comparison works for any number of reactants and any coefficients. For CH4 + 2O2 → CO2 + 2H2O, 1 mol of methane per coefficient 1 gives 1, while 3 mol of oxygen per coefficient 2 gives 1.5, so methane is limiting even though the mixture contains more oxygen molecules. The coefficient matters as much as the mole count.

Theoretical yield and leftover excess

The limiting reactant sets the product amount. Hydrogen's ratio is 0.9921, so water forms 0.9921 × 2 = 1.984 mol. With a molar mass of 18.015 g/mol, that is 1.984 × 18.015 = 35.74 g of water, which is the theoretical yield for this mixture.

The excess reactant is partly consumed. Oxygen needs 1.984 / 2 = 0.9921 mol to react with all of the hydrogen, so 1.0001 − 0.9921 = 0.0080 mol of O2 is left, about 0.256 g. The mass balance checks out: 4 g of hydrogen plus 31.74 g of reacted oxygen gives the 35.74 g of product.

Percent yield: prediction versus the lab

If the experiment collects 30 g of water, the percent yield is 30 / 35.744 × 100 = 83.9%. The theoretical yield is always based on the limiting reactant and a complete reaction, so it is the maximum possible amount for those starting amounts, not a guarantee.

Both yields must be in the same unit: two masses or two mole amounts. A percent yield below 100% is normal and can come from incomplete reaction, reversible steps, side products or losses during transfer and drying. A value above 100% usually points to a wet or impure product, an underestimated theoretical yield or a measurement error.

Assumptions behind the calculation

The arithmetic assumes an ideal, one-step reaction that goes to completion: every limiting particle reacts, no side reactions compete and no product is lost. Real systems often violate all of those assumptions a little, which is exactly why percent yield exists as a separate measurement.

It also assumes the reactants are pure and that the equation is the only process occurring. If a reactant is a hydrate or a solution, use the mass of the actual compound present, and if several reactions compete, the limiting reactant concept applies to each reaction separately.

Finally, the calculation treats every reactant as available from the start and in the same phase. In practice a reaction may be held back by mixing, temperature or catalyst availability rather than by stoichiometry, which is one reason a real yield can stay well below the theoretical value even when the arithmetic is correct.

Common mistakes and a checklist

The most common slip is comparing masses instead of moles, closely followed by dividing by the wrong coefficient or using the product coefficient instead of the reactant coefficient. Another frequent error is treating the theoretical yield as a prediction of what will actually be collected, when it is only the ceiling.

Convert every reactant to moles before comparing.
Divide by the reactant's own coefficient, not the product's.
Use the limiting reactant to scale the product amount.
Subtract consumed excess from the starting amount to find leftovers.
Convert the theoretical and actual yields to the same unit before taking the percentage.

Frequently asked questions

How do I know which reactant is limiting?

Convert each reactant amount to moles, divide each by its coefficient in the balanced equation and compare the results. The smallest quotient is the limiting reactant, as with 4 g of H2 (0.9921 after dividing by 2) versus 32 g of O2 (1.0001 after dividing by 1) in 2H2 + O2 → 2H2O.

What is theoretical yield and why does it use the limiting reactant?

Theoretical yield is the maximum product predicted by the balanced equation when the limiting reactant is completely consumed. Because the reaction stops when that reactant runs out, its coefficient ratio determines how much product can form, and the excess reactants cannot change that ceiling.

How do I calculate percent yield?

Divide the actual yield by the theoretical yield and multiply by 100. In the worked example, 30 g collected from a theoretical 35.744 g gives 30 / 35.744 × 100 = 83.9%. Both values must be masses or both must be mole amounts for the ratio to make sense.

Can percent yield be more than 100%?

It can be reported above 100%, but that does not mean atoms were created. Common causes are product that is still wet or contaminated, an underestimated theoretical yield from the wrong limiting reactant, or measurement error. Dry and purify the product before trusting a figure above 100%.

What happens to the excess reactant?

Part of it reacts in proportion to the limiting reactant and the rest is left over. In the example, 0.9921 mol of the 1.0001 mol of oxygen is consumed, leaving about 0.0080 mol, or 0.256 g, unreacted after the hydrogen is gone.

Tools for this job

All Chemistry Tools