How to Calculate Percent Yield: Limiting Reagent, Theoretical Yield and Worked Example

By Dr. Zubair Khalid, DVM, MS, PhD · 2026-09-26

Percent yield compares the amount of product you actually isolated to the amount the balanced equation says you could have made. It is the standard way to report how well a reaction performed, and it appears in nearly every organic, inorganic and analytical lab report you will write.

Getting the number right requires more than dividing two masses. You have to balance the equation, identify the limiting reagent, convert masses to moles, and decide whether the material on the filter paper is dry and pure enough to weigh. This guide walks through the full calculation, a worked aspirin example, and the assumptions that quietly break the method.

Quick Answer

Percent yield = (actual yield / theoretical yield) x 100, where theoretical yield is the mass of product expected if the limiting reagent converts completely.
The limiting reagent is the reactant that runs out first. Find it by dividing each reactant's moles by its stoichiometric coefficient; the smallest result wins.
Theoretical yield = moles of limiting reagent x (product coefficient / limiting reagent coefficient) x product molar mass.
Actual yield is what you weigh after isolation and drying, not the mass of the flask contents before workup.
A result above 100% means the sample is wet or impure, not that the reaction overperformed.

The Percent Yield Formula and What Each Term Means

The calculation rests on one ratio:

$$ \text{Percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100 $$

Each term has a precise meaning. The actual yield is the isolated mass of purified product you obtain, measured after drying. The theoretical yield is the mass of product predicted from stoichiometry, assuming the limiting reagent is consumed completely and no product is lost. Both must be in the same units, typically grams.

The theoretical yield is a ceiling, not a target. Real reactions lose material to side products, incomplete conversion, transfers between vessels, and material left on chromatography columns or filter frits. A yield of 80 to 90 percent for a clean reaction is often good; a yield of 40 percent may still be acceptable for a difficult transformation.

Because the theoretical yield depends on which reactant runs out first, you cannot compute percent yield without identifying the limiting reagent. That step comes next.

Step 1: Balance the Equation

Every stoichiometric ratio you use later comes from the balanced equation. Write it with correct formulas and whole-number coefficients. For the aspirin synthesis used below:

$$ \text{C}_7\text{H}_6\text{O}_3 + \text{C}_4\text{H}_6\text{O}_3 \rightarrow \text{C}_9\text{H}_8\text{O}_4 + \text{CH}_3\text{COOH} $$

Salicylic acid reacts with acetic anhydride in a 1:1:1 ratio to give acetylsalicylic acid (aspirin) and acetic acid. If the equation is unbalanced, the coefficient ratios will be wrong, and the limiting reagent identification will be wrong with them.

Step 2: Convert Reactant Masses to Moles

Moles connect the mass you weigh to the coefficients in the equation:

$$ n = \frac{m}{M} $$

Here $n$ is moles, $m$ is mass in grams, and $M$ is molar mass in g/mol. For a liquid reactant, you usually measure volume, so convert first using density:

$$ m = V \times \rho $$

where $V$ is volume and $\rho$ is density. Molar masses come from the periodic table or a molecular weight calculator; for hydrated salts, use the molar mass of the full hydrate, not the anhydrous form.

Step 3: Identify the Limiting Reagent

Divide each reactant's moles by its coefficient in the balanced equation:

$$ \text{comparison value} = \frac{n}{\nu} $$

where $\nu$ (nu) is the stoichiometric coefficient. The smallest comparison value identifies the limiting reagent. When all coefficients are 1, this reduces to comparing moles directly.

The limiting reagent controls the theoretical yield. Any other reactant is in excess, and some of it will remain unreacted when the reaction stops. You can calculate how much excess remains by subtracting the moles consumed (equal to the moles of limiting reagent, scaled by coefficient ratio) from the moles you started with.

Step 4: Calculate Theoretical Yield

Multiply the moles of limiting reagent by the stoichiometric ratio to product, then by the product's molar mass:

$$ \text{Theoretical yield} = n_{\text{limiting}} \times \frac{\nu_{\text{product}}}{\nu_{\text{limiting}}} \times M_{\text{product}} $$

This gives the mass of product expected under ideal conditions. It is an estimate derived from the equation, not a measurement, and it assumes complete conversion of the limiting reagent.

Worked Example: Aspirin Synthesis

The following data are illustrative. Suppose you start with 2.00 g of salicylic acid and 5.00 mL of acetic anhydride (density 1.08 g/mL), and you isolate 2.10 g of dry aspirin.

Molar masses: salicylic acid, 138.12 g/mol; acetic anhydride, 102.09 g/mol; aspirin, 180.16 g/mol.

Step 1. Moles of each reactant.

$$ n_{\text{salicylic}} = \frac{2.00\ \text{g}}{138.12\ \text{g/mol}} = 0.01448\ \text{mol} $$

$$ m_{\text{anhydride}} = 5.00\ \text{mL} \times 1.08\ \text{g/mL} = 5.40\ \text{g} $$

$$ n_{\text{anhydride}} = \frac{5.40\ \text{g}}{102.09\ \text{g/mol}} = 0.0529\ \text{mol} $$

Step 2. Identify the limiting reagent. Both coefficients are 1, so compare moles directly. Salicylic acid (0.01448 mol) is smaller than acetic anhydride (0.0529 mol), so salicylic acid is limiting.

Step 3. Theoretical yield.

$$ \text{Theoretical yield} = 0.01448\ \text{mol} \times 180.16\ \text{g/mol} = 2.609\ \text{g} $$

Step 4. Percent yield.

$$ \text{Percent yield} = \frac{2.10\ \text{g}}{2.609\ \text{g}} \times 100 = 80.5\% $$

Excess reagent remaining. Acetic anhydride consumed equals the moles of salicylic acid, 0.01448 mol. Unreacted acetic anhydride:

$$ 0.0529 - 0.01448 = 0.0384\ \text{mol} $$

which corresponds to about 3.92 g. That material stays in the reaction mixture and must be removed during workup.

If you want to check your arithmetic, the Percent Yield Calculator works out the limiting reagent, the theoretical yield and the percent yield from your reactant masses. A Moles to Grams Calculator is useful when you are converting between the two during the setup.

Quantity	Value
Salicylic acid	2.00 g, 0.01448 mol (limiting)
Acetic anhydride	5.40 g, 0.0529 mol (excess)
Theoretical yield of aspirin	2.609 g
Actual dry yield	2.10 g
Percent yield	80.5%
Unreacted acetic anhydride	0.0384 mol, about 3.92 g

Reporting Yield Honestly

A percent yield is only as good as the mass it is built on. Weigh the product after it is fully dry, and report the yield of the purified material separately from any crude yield measured before purification. A crude yield and an isolated yield answer different questions, and mixing them in a report obscures what was actually achieved.

Wernerova and Hudlicky examined the practical limits of how precisely isolated yields and stereoisomer ratios can be determined, and why reported yields deserve scrutiny [1]. Small errors in weighing, incomplete drying, and losses during purification all propagate into the final number. Treat a yield as a measured quantity with uncertainty, not an exact figure.

Percent yield also says nothing about purity. An impure sample can still give a yield close to 100 percent, because impurities add mass. Pair every yield with a purity measure such as melting point, HPLC or NMR. Yield and purity are independent pieces of evidence about a sample.

Common Mistakes

Using the wrong coefficient ratio. The product-to-limiting-reagent ratio comes straight from the balanced equation. If the equation is unbalanced, every downstream number is wrong. Fix: balance first, then read coefficients.
Forgetting to balance the equation at all. Even a simple synthesis can have non-obvious coefficients. Fix: write and check the balanced equation before touching the calculator.
Using the excess reagent to compute theoretical yield. This inflates the theoretical yield and produces a falsely low percent yield. Fix: always base the calculation on the limiting reagent.
Using a hydrate's anhydrous molar mass. If the reactant is a hydrate, the water of crystallization contributes to the mass you weighed. Fix: use the molar mass of the full hydrated formula.
Weighing product before it is fully dry. Residual solvent or water adds mass and can push the yield above 100 percent. Fix: dry to constant mass, or use a vacuum desiccator or drying oven as appropriate for the compound.
Reporting a yield above 100 percent without comment. This almost always means the product is wet or impure, retaining solvent, water, unreacted starting material or salts. Fix: re-dry and re-weigh, then check purity.

Limitations

The percent yield calculation assumes the reaction goes to completion with respect to the limiting reagent and that the only product formed is the one you want. Both assumptions can fail. Competing side reactions consume limiting reagent without producing your target, so the true conversion to product is lower than the stoichiometric prediction suggests.

The method also assumes you know the identity and purity of what you weighed. If the isolated solid is a solvate, a salt, or a mixture, the mass does not correspond to the formula you used for the molar mass, and the yield is not meaningful as calculated.

Yield is one of several metrics of reaction efficiency. Trost introduced atom economy, the fraction of reactant atoms that end up in the desired product, and a reaction with a high percent yield can still score poorly on it [2]. Constable and colleagues compared green chemistry metrics including yield, atom economy and reaction mass efficiency against the cost drivers of commercial pharmaceutical processes [3]. The US EPA and the American Chemical Society describe the principles of green chemistry, which include designing syntheses to maximize incorporation of materials into the final product [4][5]. For a full picture of a reaction's efficiency, report yield alongside at least one of these metrics.

For multi-step synthesis, the overall yield is the product of the individual step yields. Five steps at 80 percent each give $0.8^5 = 32.8\%$ overall. A modest loss at each step compounds quickly, which is why protecting-group strategy and route design matter as much as any single reaction's yield.

Frequently Asked Questions

What is the difference between actual yield and theoretical yield?

Theoretical yield is the mass of product predicted from stoichiometry if the limiting reagent converts completely. Actual yield is the mass you isolate and weigh after the reaction and workup. Percent yield is the ratio of the two, expressed as a percentage.

Why can percent yield be greater than 100 percent?

A yield above 100 percent almost always means the product is wet or impure. Retained solvent, water, unreacted starting material or salts add mass that is not your product. Dry the sample to constant mass and check purity before reporting.

Does a high percent yield mean the product is pure?

No. Percent yield says nothing about purity. A sample can contain impurities that add mass and still give a yield close to the theoretical maximum. Always pair yield with a purity measure such as melting point, HPLC or NMR.

How do I find the limiting reagent?

Convert each reactant's mass to moles, then divide each by its stoichiometric coefficient from the balanced equation. The smallest result identifies the limiting reagent. When all coefficients are 1, the reactant with the fewest moles is limiting.

How do I calculate overall yield for a multi-step synthesis?

Multiply the percent yields of each step, expressed as decimals, then convert back to a percentage. Five steps at 80 percent each give $0.8^5 = 32.8\%$ overall. Losses compound, so each step's efficiency matters.