Percent Yield Calculator

Percent yield compares how much product you actually collected with the most that the reaction could have produced. This calculator does it four ways. The simplest takes your actual and theoretical yields. The reaction mode reads a balanced equation and reactant masses, finds the limiting reagent and the theoretical yield itself, then gives the percent yield. The other two run the calculation forward: how much product to expect at a given yield, and how much starting material to use to end up with a target amount of product.

Formulas are parsed with IUPAC atomic weights, and masses can be entered in grams, kilograms or pounds. Results above 100% are flagged, since they almost always mean a wet or impure product.

percent yield = actual yield ÷ theoretical yield × 100

theoretical yield = (moles of limiting reagent ÷ its coefficient) × product coefficient × product molar mass

reactant needed = target ÷ (yield ÷ 100), converted back through the mole ratio

Theoretical yield assumes the limiting reagent reacts completely, with no side products and no losses. Actual yield is what you weigh after isolating and drying the product. Both must be in the same units and refer to the same compound. Planning backward divides the target by the expected yield fraction first, because you need to aim for more theoretical product than you hope to keep.

A simple case. Collecting 1.8 g when 2.2 g was possible is a 81.8% yield.

Aspirin. From 2.00 g of salicylic acid and excess acetic anhydride the theoretical yield is 2.61 g of aspirin. A student collecting 2.10 g has a 80.5% yield.

Iron smelting. 1 kg of Fe₂O₃ can give 699.4 g of iron; producing 600 g is 85.8%.

Planning ammonia. To collect 100 g of NH₃ from a process with 50% yield, start with 164.5 g of N₂ and excess hydrogen.

Few reactions convert every molecule of starting material. Some reach equilibrium before finishing, some form side products, and some product is always lost on glassware, filter paper and in the mother liquor during recrystallization. Teaching labs typically report 50 to 90%; optimized industrial processes can exceed 95%. A yield over 100% is a warning sign: usually the product still holds solvent or water, or contains unreacted starting material or a by-product. Drying to constant mass and checking purity, for example by melting point, usually brings it back below 100%.

A reaction can have a high percent yield and still waste most of its atoms. Atom economy measures the fraction of reactant mass that ends up in the desired product. For aspirin made from acetic anhydride, acetic acid is a by-product, so even a perfect yield leaves about a quarter of the reactant mass as waste. Green chemistry looks at both numbers, along with solvent use and energy, when comparing synthetic routes.

State the limiting reagent, the theoretical yield with its calculation, the mass of dry product collected and the percent yield, rounded to the precision of your balance, usually three significant figures. Then explain the gap: name specific losses such as product left in the flask, solubility in the wash solvent or an incomplete reaction time, rather than a general mention of human error. If the yield is above 100%, say so and give the likely cause, such as incomplete drying.