Actual Yield Calculator
Enter any two of the three values - percent yield, theoretical yield, and actual yield - to solve for the third. Switch the solve mode to find whichever quantity you need. You can work in grams or moles, and if you enter a molar mass the calculator converts between them automatically. Results update as you type and a step-by-step panel shows every calculation.
Formula
Worked example
A student reacts 15.0 g of salicylic acid with excess acetic anhydride to synthesize aspirin. Stoichiometry predicts a theoretical yield of 17.4 g. After filtering, washing, and drying, she recovers 13.9 g. Percent yield = (13.9 / 17.4) x 100 = 79.9%, a good result for this classic esterification.
What is actual yield and how does it differ from theoretical yield?
Actual yield is the mass (or moles) of product you physically collect, weigh, and record at the end of an experiment. Theoretical yield is the upper ceiling calculated from stoichiometry: it assumes the limiting reagent reacts completely and every molecule of product is isolated perfectly. In practice, actual yield is almost always below theoretical yield because some product is inevitably lost during filtration, transfer between vessels, recrystallisation, and drying. Small amounts are also lost to side reactions, incomplete reactions, and equilibrium constraints. The ratio of actual to theoretical, expressed as a percentage, is called percent yield, and it is the standard measure of how efficiently a reaction was carried out.
How to calculate actual yield step by step
Start by identifying the limiting reagent - the reactant that is completely consumed first and caps the amount of product that can form. Use the molar mass of the limiting reagent to convert its mass to moles, then use the mole ratio from the balanced equation to find the theoretical yield in moles of product, and finally convert back to grams using the product molar mass. That is the theoretical yield. If you already know the percent yield from previous runs or literature, multiply: actual yield = (percent yield / 100) x theoretical yield. For example, if the theoretical yield is 9.00 g and the percent yield is 80.0%, the actual yield is (80.0 / 100) x 9.00 = 7.20 g. If instead you have the actual and theoretical yields, divide actual by theoretical and multiply by 100 to get percent yield. This calculator handles all three arrangements so you only need to supply the two values you know.
Why percent yield is almost never 100%
A true 100% yield is physically unachievable because some product is always mechanically lost during the isolation process. Even a small drop left on glassware, a few crystals that pass through a filter, or a thin film on a spatula reduces the measured yield. Beyond mechanical losses, reactions can reach chemical equilibrium before all the limiting reagent is consumed, side reactions divert some material to unwanted products, and reagents or products may decompose under the reaction conditions. In teaching laboratories a yield of 70% or higher is generally considered satisfactory. Industrial processes are designed to maximise yield because even a 1% improvement across millions of kilograms of product represents significant savings. A percent yield above 100% is a red flag: it indicates the product was not fully dried, a solvent or reagent is still incorporated, or there was a calculation error in the theoretical yield.
Units, molar mass conversion, and comparing runs
Both actual and theoretical yield must be in the same units for the formula to work, whether that is grams, moles, kilograms, or any other consistent measure. If your experiment reports in moles but your literature value is in grams, use the product molar mass to convert: mass (g) = moles x molar mass (g/mol), and moles = mass (g) / molar mass (g/mol). This calculator accepts grams, moles, kilograms, and milligrams, and if you enter the molar mass it will show the converted result automatically. When comparing yields across different experiments, always check that you are using the same scale: a yield of 0.025 mol is not directly comparable to 4.5 g unless you convert both to the same unit first.
Typical percent yield ranges by context
| Percent yield | Rating | Common interpretation |
|---|---|---|
| > 100% | Invalid | Product contains impurities or is not fully dried |
| 90-100% | Excellent | Minimal loss; high-purity isolation technique |
| 70-89% | Good | Acceptable for most teaching and research labs |
| 50-69% | Moderate | Significant losses; review conditions and purification |
| < 50% | Low | Poor recovery; possible side reactions or incorrect limiting reagent |
General benchmarks used in teaching labs, industrial synthesis, and research to evaluate reaction efficiency.
Frequently asked questions
What is the formula for actual yield?
Actual yield = (percent yield / 100) x theoretical yield. Rearranging the standard percent yield formula gives this: if percent yield = (actual / theoretical) x 100, then actual = theoretical x (percent yield / 100). For example, a theoretical yield of 9.00 g at 80.0% gives an actual yield of 7.20 g.
Can actual yield be greater than theoretical yield?
Not in a physically meaningful way. If your measured actual yield exceeds the theoretical yield, the percent yield will be above 100%, which is impossible for a pure product. The most common reason is that the product is not fully dry and contains absorbed water or solvent. Return the product to the desiccator, dry to constant mass, and reweigh. Less commonly, the theoretical yield was calculated incorrectly, such as using the wrong mole ratio or the wrong limiting reagent.
What is percent yield and why does it matter?
Percent yield = (actual yield / theoretical yield) x 100. It measures how efficiently a reaction was performed. A high percent yield means little product was lost during isolation and purification. In research it is a benchmark of technique, and in industry it directly affects cost because wasted starting material and product translate into direct financial losses. Most teaching laboratories consider 70% or higher a good result for a first attempt.
How do I find the theoretical yield?
Write and balance the chemical equation. Identify the limiting reagent (the one that runs out first). Convert the mass of the limiting reagent to moles using its molar mass. Multiply by the mole ratio (moles of product per mole of limiting reagent from the balanced equation). Convert back to grams using the product molar mass. Alternatively, if you know both the actual yield and the percent yield, this calculator rearranges the formula for you: theoretical yield = actual yield / (percent yield / 100).
What is a limiting reagent and how does it relate to yield?
The limiting reagent (or limiting reactant) is the reactant that is completely consumed during the reaction, stopping further product from forming. It sets the maximum possible yield. The other reagents present in excess do not constrain the yield. Identifying the wrong limiting reagent is a common calculation error that produces an incorrect theoretical yield, which in turn makes the percent yield meaningless. To find the limiting reagent, convert all reactant masses to moles and compare the mole ratios to the stoichiometric coefficients.
What are common reasons for a low percent yield?
Typical reasons include: product lost on glassware during transfer; incomplete reaction because equilibrium was reached before all limiting reagent was consumed; competing side reactions that convert some starting material to unwanted products; product decomposed during the reaction or workup; poor technique during recrystallisation, filtration, or distillation; and evaporation of volatile products. Improving yield usually involves optimising reaction time, temperature, solvent, and purification procedure.