What is Actual Yield? Actual Yield Definition
The actual yield is the quantity of a product actually obtained from a chemical reaction in the laboratory. No chemical reaction is ever carried out with perfect efficiency — losses occur through incomplete reactions, side products, transfer losses, and measurement errors. Because of this, the actual yield is almost always lower than the theoretical yield (the maximum amount predicted by stoichiometry).
Actual Yield Formula
The actual yield is calculated using the following formula:
Ya = (Yp / 100) × Yt
Where:
- Ya — Actual yield (in grams, kg, mg, oz, or lb)
- Yp — Percent yield (%)
- Yt — Theoretical yield (same mass unit as Ya)
Difference Between Theoretical Yield and Actual Yield
The theoretical yield is calculated from the balanced chemical equation and assumes perfect, complete conversion of the limiting reactant into product. The actual yield is what you measure after the reaction is done. The difference arises because:
- Reactions rarely go to 100% completion.
- Some product may remain on glassware, filter paper, or in solution.
- Side reactions consume some reactants.
- Product may decompose or evaporate under reaction conditions.
In rare cases the actual yield can appear to exceed the theoretical yield — this indicates impurities (residual solvent, unreacted starting material) or a weighing error.
Units Supported
This calculator supports both metric and imperial (American) units of mass:
- Metric: grams (g), kilograms (kg), milligrams (mg)
- Imperial: ounces (oz), pounds (lb)
Examples of How to Find Actual Yield
Example 1
A reaction has a percent yield of 45% and a theoretical yield of 4 g. What is the actual yield?
Ya = (45 / 100) × 4 g = 1.8 g
Example 2
The decomposition of magnesium carbonate (MgCO3 → MgO + CO2) has an experimental percent yield of 79% and a theoretical yield of 19 g. What is the actual yield of magnesium oxide?
Ya = (79 / 100) × 19 g = 15.01 g
Frequently Asked Questions
Why is actual yield usually less than theoretical yield?
Because real reactions have incomplete conversion, side reactions, transfer losses, and solubility effects. A reaction that proceeds to 100% completion with zero loss is essentially impossible in practice.
Can actual yield be greater than theoretical yield?
Not in a perfectly pure system. If the measured actual yield exceeds the theoretical yield, the product contains impurities (unreacted reagents, solvent) or there is a measurement error.
How do I find the theoretical yield first?
Use a theoretical yield calculator. You need the balanced equation, the masses of reactants, and their molar masses. The limiting reactant determines the theoretical yield.
How is percent yield related to actual yield?
Percent yield = (actual yield / theoretical yield) × 100%. Use our percent yield calculator to solve the reverse problem.