When chemists perform a reaction, they combine starting materials, called reactants, to create new substances, or products. Quantifying the amount of product successfully made determines the efficiency of the chemical process. This measured quantity of the final substance obtained from a completed experiment is known as the actual yield.
Defining Actual Yield
The actual yield is the measured mass of the desired product collected at the end of a chemical reaction. It represents the real-world outcome, reflecting the amount of material physically isolated and weighed in the laboratory. This value is empirical, meaning it is derived solely through physical measurement after the reaction is performed, not through calculation beforehand. The actual yield is typically expressed in mass units, such as grams, or sometimes in moles.
Actual Versus Theoretical Yield
To understand the practical meaning of the actual yield, it must be compared to the theoretical yield. The theoretical yield is the maximum possible amount of product that could be formed from the initial amounts of reactants, assuming the reaction proceeds perfectly. This value is calculated using the balanced chemical equation and stoichiometry, which describes the mole ratios between reactants and products. The theoretical yield is determined by the limiting reactant, which is the reactant that will be completely consumed first. The actual yield is what is obtained after the experiment, while theoretical yield is what should be obtained under ideal conditions.
Quantifying Efficiency: Calculating Percent Yield
Chemists use both the actual yield and the theoretical yield to determine the efficiency of a reaction through a metric called percent yield. Percent yield expresses the actual amount of product obtained as a percentage of the maximum possible amount. The formula for this measurement is: Percent Yield = (Actual Yield / Theoretical Yield) \(\times\) 100%.
Interpreting the percent yield provides a clear measure of the reaction’s success in a laboratory or industrial setting. For example, a percent yield of 90% indicates that the experiment successfully converted 90% of the limiting reactant into the desired product. This calculation is a fundamental tool for evaluating and optimizing chemical processes, as a higher percentage signifies a more efficient use of starting materials.
Why Actual Yield Is Rarely Perfect
The actual yield is almost always lower than the calculated theoretical yield, meaning the percent yield is usually less than 100%. One primary reason is that many reactions are reversible and reach a chemical equilibrium, which prevents the complete conversion of all reactants into products. This incomplete reaction limits the final amount of product that can be formed.
Another common factor is the occurrence of side reactions, where reactants interact in unintended ways to produce unwanted byproducts. These competing reactions consume starting materials intended for the desired product, lowering the final actual yield. Furthermore, product is invariably lost during handling and purification, such as when transferring substances or filtering solids. Some product might stick to the glassware or dissolve in the wash solvent, leading to a measured mass less than the theoretical maximum.