Find the percent yield of a reaction, or solve for the actual or theoretical yield. The theoretical yield helper works it out from the limiting reactant, with every step shown.
The percent yield formula
Percent yield = actual yield ÷ theoretical yield × 100
The theoretical yield is the most product a reaction can make if all of the limiting reactant reacts and nothing is lost. The actual yield is what you really collect and weigh. Use the same unit for both, usually grams.
The theoretical yield formula
Theoretical yield (g) = moles of limiting reactant × (product coefficient ÷ reactant coefficient) × molar mass of product
- Balance the equation.
- Find the limiting reactant: the one that runs out first once you account for the mole ratio.
- Convert its mass to moles: mass ÷ molar mass.
- Use the coefficients to get moles of product.
- Multiply by the product's molar mass to get grams.
Worked example
Procaine (Novocain) is made from p-aminobenzoic acid (137.14 g/mol) and 2-diethylaminoethanol in a 1:1 ratio. With 10.0 g of each, the acid is limiting: 10.0 ÷ 137.14 = 0.07292 mol. One mole of acid gives one mole of procaine (236.31 g/mol), so the theoretical yield is 0.07292 × 236.31 = 17.23 g. If you collect 15.7 g, the percent yield is 15.7 ÷ 17.23 × 100 = 91.11%. (Rounding the theoretical yield to 17.2 g first gives 91.28%, so keep extra digits until the end.)
Why yields are below 100%, and what over 100% means
Real reactions lose product: some reactants don't react, side reactions make other products, and some product stays behind on glassware or in solution during purification. A yield over 100% almost always means the product is still wet with solvent or impure, or a mass was weighed wrong.
Limitations
- You need to know which reactant is limiting. If you're not sure, work out the theoretical yield from each reactant; the smaller answer is the real one.
- The helper assumes one product from one limiting reactant, as written in the balanced equation.
- Need a molar mass? Use the molar mass calculator. Comparing a measurement to a known value instead? That's percent error.
Sources
How we checked this
Before every update, this calculator is run in a real browser with these known answers, and it must match all of them:
- 15.7 g actual ÷ 17.2 g theoretical = 91.28% ✓
- 10.0 g of a 137.14 g/mol reactant, 1:1, product 236.31 g/mol: 17.2313 g theoretical ✓
- …and 15.7 g collected is a 91.11% yield ✓
- 2 mol H₂ with 2H₂ + O₂ → 2H₂O (18.015 g/mol): 36.03 g of water ✓
- 85% of a 50 g theoretical yield is 42.5 g ✓
- Edge case: a theoretical yield of 0 asks for a value above 0 ✓
- Over 100% (20 g from 17.2 g) = 116.28% and a warning ✓
Last checked October 26, 2026. Spotted something wrong? The math is done in your browser with the formula shown on this page, so you can check any result by hand.
Questions people ask
What is the percent yield formula?
Percent yield = actual yield ÷ theoretical yield × 100. 15.7 g collected from a theoretical 17.2 g is 91.28%.
How do you calculate theoretical yield?
Convert the limiting reactant to moles, multiply by the mole ratio from the balanced equation (product ÷ reactant coefficient), then multiply by the product's molar mass.
How do you calculate percent yield from grams?
Divide the grams you collected by the theoretical grams and multiply by 100. Both must be in the same unit.
Can percent yield be over 100%?
Not for a pure, dry product. Over 100% usually means leftover solvent, impurities or a weighing mistake.
What is the difference between percent yield and percent error?
Percent yield compares what you made with the most you could make. Percent error compares a measurement with an accepted value and is always about accuracy.