Percent Yield

Percent yield = (actual yield ÷ theoretical yield) × 100%: the fraction of the predicted product a reaction really delivers.

\%\text{ yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%

percent yield = (actual yield / theoretical yield) x 100%

What it means

The theoretical yield is what perfect stoichiometry promises from the limiting reactant; the actual yield is what you really weigh at the end. Their ratio, expressed as a percent, grades the reaction's efficiency. Anything below 100% reflects the real world: side reactions, incomplete conversion, and product left behind on glassware and filter paper.

Variables

Symbol	Name	Unit	Meaning
$actual yield$	Actual yield	g or mol	The mass or moles of product genuinely obtained and isolated from the reaction.
$theoretical yield$	Theoretical yield	g or mol	The maximum mass or moles of product the stoichiometry predicts from the limiting reactant.
$% yield$	Percent yield	%	The efficiency of the reaction as a percentage; always 100% or less for a clean experiment.

Assumptions

Both yields are expressed in the same units — grams with grams, moles with moles.
The theoretical yield comes from the limiting-reactant stoichiometry.
A result above 100% signals measurement error or an impure, contaminated product.
The actual yield is the mass or moles genuinely isolated, not estimated.

Worked example

Problem

A synthesis has a theoretical yield of 25.0 g, but only 18.4 g of pure product is isolated. Calculate the percent yield.

Given:

Theoretical yield = 25.0 g
Actual yield = 18.4 g

Write % yield = (actual yield / theoretical yield) × 100%.
Substitute: % yield = (18.4 g / 25.0 g) × 100%.
Divide: 18.4 / 25.0 = 0.736.
Multiply by 100%: 0.736 × 100% = 73.6%.

73.6%

The reaction delivered about three-quarters of the stoichiometric promise — respectable, with the missing 26.4% lost to side reactions and handling.