Every reaction has a best-case ceiling — the theoretical yield you would get if the chemistry went perfectly and you lost nothing. Reality always falls short: side reactions, incomplete conversion and losses in handling mean the actual yield is smaller. The percent yield measures how close you came, as a simple ratio times 100. It is the headline number of any synthesis — a measure of how efficient the reaction and your technique really were, and the figure every lab report and process chemist lives by.
Reviewed: June 19, 2026 · Author: Naveen P N, Founder — AI Calculator · Verified against: the definition of percentage yield.
The percent-yield equations
Both yields must be measured in the same units — grams with grams, or moles with moles — because the ratio relies on like-for-like quantities; the units then cancel and leave a pure percentage. The theoretical yield comes from a stoichiometry calculation on the limiting reactant, while the actual yield is whatever you weigh out, dried and pure. The gap between them is the product lost along the way.
Worked example — a synthesis
Scenario: A reaction has a theoretical yield of 10 g, and you isolate 8.5 g of dry pure product.
An 85% yield is a good result — 1.5 g, or 15%, was lost to side reactions and handling. If you only recovered 5 g, the yield would drop to 50%, signalling something to improve in the method. And if you weighed 10.5 g, the apparent 105% yield would be a red flag that the product is still wet with solvent or contaminated, not that you broke conservation of mass.
Frequently Asked Questions
(actual ÷ theoretical) × 100. 8.5 g actual from a 10 g theoretical = 85%. Same units both sides.
Theoretical = max from the balanced equation/limiting reactant; actual = what you really isolate. Actual ≤ theoretical.
Side reactions, incomplete reaction, losses in filtering/transfer, impure reagents. 70–90% is typically good.
No — it signals wet/impure product or a wrong theoretical value. Dry fully and recheck the limiting reactant.
From the balanced equation: limiting reactant → moles → mole ratio → moles product → grams via molar mass.