Free Percent Yield Calculator
Calculate the percent yield of a chemical reaction by dividing the actual mass you recovered by the theoretical mass and multiplying by 100, giving an instant measure of how efficient your reaction or process was.
Enter the actual and theoretical yields in the same unit (g, mg, or mol) — percent yield is calculated live.
% yield = (actual yield ÷ theoretical yield) × 100. The theoretical yield comes from the limiting reactant via the balanced equation. Keep both values in the same unit so they cancel to a clean percentage.
Quick answer
Percent yield = (actual yield ÷ theoretical yield) × 100. For example, if a reaction produces 4.5 g of product when the stoichiometry predicts 5 g, the percent yield is (4.5 ÷ 5) × 100 = 90%. Use the same mass unit (g, mg, mol) for both values so the units cancel.
Formula & method
% yield = (actual yield ÷ theoretical yield) × 100
- actual yield — Mass (or moles) of product actually recovered and measured in the lab
- theoretical yield — Maximum mass (or moles) predicted by stoichiometry from the limiting reactant
Both yields must be in the same unit (grams, milligrams, or moles) so the ratio is dimensionless. The result is a percentage, normally between 0% and 100%.
actual yield = (% yield ÷ 100) × theoretical yield
Rearranged form used to predict how much product you should expect when a typical percent yield for the reaction is known.
Examples
- Input
- Actual = 4.5 g, Theoretical = 5 g
- Result
- 90%
- Why
- (4.5 ÷ 5) × 100 = 0.9 × 100 = 90%. The reaction recovered 90% of the maximum possible product, a strong result for a single-step synthesis.
- Input
- Actual = 27.6 g, Theoretical = 35 g
- Result
- 78.8571%
- Why
- (27.6 ÷ 35) × 100 = 0.788571… × 100 = 78.8571%. About 21% of the theoretical product was lost to side reactions, incomplete conversion, or transfer/filtration losses.
- Input
- Actual = 8.92 g, Theoretical = 10.0 g
- Result
- 89.2%
- Why
- (8.92 ÷ 10.0) × 100 = 0.892 × 100 = 89.2%. Working in milligrams (8920 mg ÷ 10000 mg) gives the identical 89.2% because the units cancel.
When to use this tool
- Reporting the efficiency of a synthesis or precipitation in a chemistry lab write-up.
- Comparing two reaction routes or conditions to see which actually delivers more product.
- Estimating how much starting material you need to hit a target amount of product, using a known typical yield.
- Checking homework or exam problems that ask for percent yield from actual and theoretical masses.
Common mistakes
- Swapping the inputs — dividing theoretical by actual gives a number above 100% and inverts the meaning. Actual goes on top, theoretical on the bottom.
- Mixing units between the two boxes, such as grams for actual and moles for theoretical. Convert both to the same unit first so the ratio is valid.
- Forgetting the theoretical yield must come from the limiting reactant. Using the excess reagent's amount inflates the theoretical value and understates your real efficiency.
- Treating a percent yield over 100% as success. It almost always signals impure, wet, or unreacted product still weighed in with the desired compound.
Frequently asked questions
What is the formula for percent yield?
Percent yield equals the actual yield divided by the theoretical yield, multiplied by 100. The actual yield is what you physically recover; the theoretical yield is the maximum predicted by stoichiometry from the limiting reactant. Use the same unit for both so the answer is a clean percentage.
Can percent yield be more than 100%?
In principle no, because you cannot create more product than the reactants allow. A value above 100% is a red flag that the measured product still contains impurities, leftover solvent or water, or unreacted starting material. Dry, purify, and reweigh before recalculating.
What is the difference between actual and theoretical yield?
Theoretical yield is the ideal amount of product calculated from the balanced equation assuming the limiting reactant converts completely with no losses. Actual yield is the smaller, real amount you isolate after the reaction, side reactions, and workup. Percent yield compares the two.
Why is my percent yield so low?
Common causes are incomplete reactions, competing side reactions, product lost during filtration or transfer, an impure or wet limiting reactant, or product still dissolved in the mother liquor. Low yields are normal for multi-step syntheses, where each step multiplies the losses.
Do the actual and theoretical yields need the same units?
Yes. Both must be in the same unit — grams, milligrams, or moles — so the units cancel and you get a pure ratio. Mixing grams with moles produces a meaningless number. If your data is in moles for one and grams for the other, convert first using molar mass.
How do I find the theoretical yield first?
Identify the limiting reactant, convert its mass to moles, apply the mole ratio from the balanced equation to find moles of product, then convert back to grams using the product's molar mass. That gram figure is the theoretical yield you enter into the calculator.
Sources & references
External references open in a new tab. We are independent and not affiliated with these organizations.
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Provided “as is” for general information only — results may be inaccurate, so verify before you rely on them. No warranty; use at your own risk.
Built and reviewed by HIFreeTools against the formula shown above and any authoritative references cited on this page. See our methodology and editorial standards.
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