Calculate percent yield from actual and theoretical yield, or find theoretical yield from the limiting reactant. Solve for any value with steps.
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Percent yield measures how efficient a chemical reaction was — the amount of product you actually obtained compared to the maximum the reaction could produce in theory. It's calculated as the actual yield divided by the theoretical yield, times 100. This calculator works both ways: enter your actual and theoretical yields to get the percent, or give it any two of the three values (percent, actual, theoretical) and it solves for the third. It can also compute the theoretical yield for you from the limiting reactant using stoichiometry — just enter the reactant mass, the molar masses, and the mole ratio from your balanced equation. You'll get the result with a quality rating and the full working, no signup required.
Percent yield = (actual yield ÷ theoretical yield) × 100. The actual yield is what you measured in the lab; the theoretical yield is the maximum possible, calculated from the limiting reactant. For example, if you obtained 14.9 g of product and the theoretical maximum was 15.7 g, the percent yield is (14.9 ÷ 15.7) × 100 = 94.9%. To find the theoretical yield from the limiting reactant: convert its mass to moles (mass ÷ molar mass), multiply by the mole ratio of product to reactant from the balanced equation, then multiply by the product's molar mass. So 10 g of a reactant with molar mass 40 g/mol in a 1:1 reaction making a product of molar mass 58.44 g/mol gives 0.25 mol × 58.44 = 14.61 g theoretical. As a rough guide, a yield of 90%+ is excellent, 80–89% very good, 70–79% good, and below 40% is poor. Percent yield is normally at or below 100% — a value above 100% signals impurities, leftover solvent, or measurement error.
Percent Yield Formula
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Find percent yield, or work backwards to the actual or theoretical yield.
Compute the theoretical yield from the limiting reactant with stoichiometry built in.
See instantly whether your yield is excellent, good, or poor.
Flags yields above 100% that point to impurities or measurement mistakes.
Shows the formula and each step — ideal for homework and lab reports.
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Divide the actual yield (what you obtained) by the theoretical yield (the maximum possible), then multiply by 100. For example, 14.9 g obtained out of a 15.7 g theoretical maximum is (14.9 ÷ 15.7) × 100 = 94.9%.
Theoretical yield is the maximum amount of product a reaction can make, set by the limiting reactant. Convert the limiting reactant's mass to moles, multiply by the product-to-reactant mole ratio from the balanced equation, then multiply by the product's molar mass. This calculator's theoretical-yield mode does all of that.
Actual yield is the amount of product you actually obtain and measure from a reaction. It's always less than the theoretical yield in practice, because of side reactions, incomplete reactions, and losses during transfer and purification.
It depends on the reaction, but as a general guide: 90% or above is excellent, 80–89% is very good, 70–79% is good, 40–69% is moderate, and below 40% is poor. Complex, multi-step syntheses often have much lower yields than simple reactions.
Not in theory — you can't make more product than the reactants allow. A calculated value above 100% almost always means the product still contains impurities or solvent (it wasn't fully dried), or there was a weighing or calculation error.
Percent yield = (actual yield ÷ theoretical yield) × 100%. You can rearrange it to find actual yield (theoretical × percent ÷ 100) or theoretical yield (actual ÷ percent × 100).
Convert each reactant's mass to moles, then divide by its coefficient in the balanced equation. The reactant with the smallest result is the limiting reactant — it runs out first and sets the theoretical yield.
Real reactions lose product to competing side reactions, reactions that don't go fully to completion, reversible equilibria, and material lost during filtering, transferring, and purifying. That's why actual yield is almost always below the theoretical maximum.