When chemists perform a chemical reaction, their goal is to transform starting materials into a desired product. To evaluate the success of this process, they must compare the maximum possible quantity of product that should be made (theoretical yield) with the amount that is actually produced (actual yield). This comparison relies on two fundamental measurements: the theoretical yield and the actual yield. Understanding the difference between these two values determines the efficiency and practicality of any chemical synthesis.
Defining Theoretical Yield
The theoretical yield represents the maximum mass of product a chemical reaction can generate under ideal conditions. This value is calculated mathematically, not measured in a laboratory. The calculation relies entirely on the balanced chemical equation, which provides the precise molecular ratios (stoichiometry) between the reactants and the product.
To determine this maximum output, chemists must first identify the limiting reactant. This is the starting material that will be completely consumed first, setting a ceiling for the amount of product that can form. The theoretical yield assumes the reaction proceeds with perfect efficiency, converting every molecule of the limiting reactant into the desired product with no losses.
Defining Actual Yield
In contrast to the calculated maximum, the actual yield is the quantity of product measured after a chemical reaction has been performed in a laboratory setting. This experimentally derived result is typically determined by weighing the purified and dried product. The actual yield reflects the practical outcome of the synthesis.
Obtaining this value involves steps like filtration, drying, and purification, which are necessary to isolate the desired compound. Because of inherent challenges in transferring and purifying materials, the measured actual yield is nearly always lower than the calculated theoretical yield.
Calculating the Reaction’s Efficiency (Percent Yield)
The relationship between the theoretical and actual yields is quantified by the percent yield, which evaluates a reaction’s efficiency. The percent yield is calculated by dividing the actual yield by the theoretical yield and multiplying the result by 100.
For example, if a reaction was calculated to produce a theoretical yield of 10 grams but yielded 8 grams of product, the percent yield is 80%. This indicates the reaction was 80% efficient at converting the starting material into the desired product. A high percentage signifies a highly efficient process, which is important in industrial production to minimize waste and cost.
Why Actual Yield Falls Short
The actual yield is less than the theoretical yield because real-world chemical reactions rarely proceed perfectly. A common source of this difference is an incomplete reaction, where the reactants fail to convert entirely into products, often due to reaching chemical equilibrium. This leaves a portion of the starting material unreacted, lowering the final product mass.
Another factor is the occurrence of side reactions, where starting materials react in unintended ways to produce byproducts instead of the desired compound. These competing reactions consume the reactants, diverting them away from the main pathway.
Product loss during the isolation and purification steps also contributes significantly to a lower actual yield. These losses occur during procedures such as:
- Transferring liquids between containers, where residual material is often left behind.
- Filtering solids, where some product may remain on the filter paper or glassware.
- Drying the final material, which can sometimes lead to sublimation or decomposition.
- Impurity: If the recovered product is not perfectly pure, containing trace amounts of solvent or unreacted starting material, the measured mass may be inaccurately high, though the true yield of the desired compound is still reduced.
These experimental limitations ensure that a 100% percent yield is an expectation only in theory.