Percent Yield Calculator
Calculate the percent yield of a reaction from the actual and theoretical yield. If you don't know the theoretical yield yet, work it out from the limiting reactant's mass, the balanced equation's coefficients and the chemical formulas (molar masses are calculated for you). A third mode tells you how much reactant to start with.
Optional when finding the theoretical yield
From the balanced equation
How to use the Percent Yield Calculator
- Choose a mode: percent yield, theoretical yield from a reaction, or reactant needed.
- For percent yield, enter the actual and theoretical yields in the same unit.
- For a reaction, enter the limiting reactant mass, both formulas (or molar masses) and the coefficients from the balanced equation.
- Optionally add the actual yield to get the percent yield in the same step.
- Read the result and the step-by-step stoichiometry.
Percent yield and theoretical yield formulas
- Percent yield = actual yield ÷ theoretical yield × 100%
- Moles of reactant = mass ÷ molar mass
- Moles of product = moles of limiting reactant × (product coefficient ÷ reactant coefficient)
- Theoretical yield = moles of product × product molar mass
- Reactant needed = (wanted mass ÷ expected yield) ÷ Mproduct × (reactant coefficient ÷ product coefficient) × Mreactant
Worked examples
| Reaction | Start | Theoretical yield | Actual | % yield |
|---|---|---|---|---|
| C₇H₆O₃ → C₉H₈O₄ (aspirin), 1 : 1 | 2.00 g salicylic acid (138.12 g/mol) | 2.609 g | 2.10 g | 80.5% |
| CaCO₃ → CaO + CO₂, 1 : 1 | 10.0 g CaCO₃ (100.09 g/mol) | 5.603 g CaO | 5.10 g | 91.0% |
| 2H₂ + O₂ → 2H₂O, 2 : 2 | 4.00 g H₂ (2.016 g/mol) | 35.74 g H₂O | 32.0 g | 89.5% |
Molar masses are computed from IUPAC standard atomic weights. In the formula boxes you can use brackets and hydrate dots, for example Ca(OH)2, K4[Fe(CN)6] or CuSO4·5H2O, or simply type the molar mass as a number.
Frequently asked questions
What is the percent yield formula?
Percent yield = (actual yield ÷ theoretical yield) × 100%. The actual yield is the mass of product you really obtained; the theoretical yield is the maximum mass the balanced equation allows from the limiting reactant. Both must be in the same unit.
How do I calculate theoretical yield?
Convert the limiting reactant's mass to moles (mass ÷ molar mass), multiply by the mole ratio from the balanced equation (product coefficient ÷ reactant coefficient), then convert the moles of product back to grams (moles × product molar mass).
Worked example: aspirin from salicylic acid
C7H6O3 + C4H6O3 → C9H8O4 + C2H4O2. 2.00 g of salicylic acid (138.12 g/mol) is 0.01448 mol, giving at most 0.01448 mol of aspirin (180.16 g/mol) = 2.609 g. If you collect 2.10 g, the percent yield is 2.10 ÷ 2.609 × 100 ≈ 80.5%.
Can percent yield be over 100%?
Not for a pure, dry product. A yield above 100% usually means the product is still wet with solvent, contains impurities or unreacted starting material, or there was a weighing or calculation error.
Why is the actual yield usually lower than the theoretical yield?
Reactions may not go to completion, side reactions form other products, and some product is lost during filtering, transferring, washing or recrystallisation. Yields of 70–90% are common for a single well-run laboratory step.
What is the limiting reactant?
It is the reactant that runs out first, so it determines the maximum amount of product. To find it, divide each reactant's moles by its coefficient in the balanced equation; the smallest value is the limiting reactant.
What is the difference between percent yield and percent error?
Percent yield compares what you obtained with the theoretical maximum. Percent error measures how far a measured value is from an accepted value: |measured − accepted| ÷ accepted × 100.