The percent yield is a comparison between the actual yield—which is the weight of the intended product of a chemical reaction in a laboratory setting—and the theoretical yield—the measurement of pure intended isolated product, based on the chemical equation of a flawless chemical reaction, and is defined as,
The  ideal relationship between products and reactants in a chemical reaction can be obtained by using a chemical reaction equation. Stoichiometry is used to run calculations about chemical reactions, for example, the stoichiometric mole ratio between reactants and products. The stoichiometry of a chemical reaction is based on chemical formulas and equations that provide the quantitative relation between the number of moles of various products and reactants, including yields. Stoichiometric equations are used to determine the limiting reagent or reactant—the reactant that is completely consumed in a reaction.  The limiting reagent determines the theoretical yield—the relative quantity of moles of reactants and the product formed in a chemical reaction. Other reactants are said to be present in excess. The actual yield—the quantity physically obtained from a chemical reaction conducted in a laboratory—is often less than the theoretical yield. The theoretical yield is what would be obtained if all of the limiting reagent reacted to give the product in question. A more accurate yield is measured based on how much product was actually produced versus how much could be produced. The ratio of the theoretical yield and the actual yield results in a percent yield.
When more than one reactant participates in a reaction, the yield is usually calculated based on the amount of the limiting reactant, whose amount is less than stoichiometrically equivalent (or just equivalent) to the amounts of all other reactants present. Other reagents present in amounts greater than required to react with all the limiting reagent present are considered excess. As a result, the yield should not be automatically taken as a measure for reaction efficiency.
In their 1992 publication General Chemistry, Whitten, Gailey, and Davis described the theoretical yield as the amount predicted by a stoichiometric calculation based on the number of moles of all reactants present. This calculation assumes that only one reaction occurs and that the limiting reactant reacts completely.
According to Whitten, the actual yield is always smaller (the percent yield is less than 100%), often very much so, for several reasons. As a result, many reactions are incomplete and the reactants are not completely converted to products. If a reverse reaction occurs, the final state contains both reactants and products in a state of chemical equilibrium. Two or more reactions may occur simultaneously, so that some reactant is converted to undesired side products. Losses occur in the separation and purification of the desired product from the reaction mixture. Impurities are present in the starting material which do not react to give desired product.