Percent Yield Calculator
Percent yield from the actual and theoretical yield, or the theoretical yield itself from the limiting reactant’s mass, the molar masses and the mole ratio.
Result…
How it's calculated
Percent yield = actual yield ÷ theoretical yield × 100 Theoretical yield = (reactant mass ÷ reactant molar mass) × (product coefficient ÷ reactant coefficient) × product molar mass
The theoretical yield is the most product the limiting reactant can make. Convert its mass to moles with its molar mass, scale by the mole ratio from the balanced equation (product coefficient over reactant coefficient), then convert back to mass with the product’s molar mass. The percent yield compares what you actually collected with that maximum. Both yields must be in the same unit; this calculator keeps them in the one you choose. Mass units convert exactly (1 lb = 453.59237 g, 1 oz = 28.349523125 g).
Your numbers, step by step
About the Percent Yield calculator
This percent yield calculator compares the product you actually collected with the most the reaction could have made. If you already know the theoretical yield, enter it with the actual yield and it gives the percent yield, the mass not recovered and the percentage lost. If you don't, enter the mass of the limiting reactant, its molar mass, the product's molar mass and their coefficients from the balanced equation: it works out the theoretical yield through moles, then the percent yield. Students use it for lab reports, and anyone scaling up a synthesis uses it to judge how much material is lost in each step. Masses can be in micrograms, milligrams, grams, kilograms, ounces or pounds, and the steps show the mole conversion and ratio, so you can check each one.
How to use it
- Under “I know”, choose “Actual and theoretical yield” or “The limiting reactant”.
- Choose the “Mass unit” and enter the “Actual yield”: the dry mass of product you weighed.
- With the first option, enter the “Theoretical yield” in the same unit.
- With the second, enter the “Mass of the limiting reactant”, the “Molar mass of the reactant”, the “Molar mass of the product” and the “Coefficient of the reactant” and “Coefficient of the product” from the balanced equation.
- Read the theoretical yield, the percent yield and how much product was not recovered.
Worked example
Making aspirin (C9H8O4, 180.159 g/mol) from 10 g of salicylic acid (C7H6O3, 138.121 g/mol) in a 1 : 1 ratio: 10 ÷ 138.121 = 0.0724 mol of reactant, so 0.0724 mol of aspirin, × 180.159 = 13.04 g theoretical yield. Collecting 9.5 g gives 9.5 ÷ 13.04 × 100 = 72.83%, with 3.54 g not recovered.
Updated
Questions people ask
How do I calculate percent yield?
Divide the actual yield by the theoretical yield, both in the same unit, and multiply by 100. 9.5 g collected out of a possible 12 g is 79.17%.
How do I work out the theoretical yield from grams of reactant?
Divide the limiting reactant's mass by its molar mass, multiply by the product's coefficient over the reactant's coefficient, then multiply by the product's molar mass. The stoichiometry calculator finds which reactant is limiting first.
Is a percent yield over 100% possible?
Not for a pure, dry product. A result above 100% usually means the product still holds solvent or water, has impurities, or the theoretical yield was based on the wrong reactant.
What mistake gives a wrong theoretical yield most often?
Using the reactant that is in excess instead of the limiting one, or forgetting the mole ratio when the coefficients are not 1 : 1.
Why is my yield always lower than the theoretical yield?
Real reactions rarely convert everything: some product is lost on filter paper and glassware, some reactant forms side products, and some reactions stop at an equilibrium.