The actual yield of a chemical reaction often falls short of the theoretical yield due to several factors, including incomplete reactions, side reactions, and product loss during purification.

In an ideal scenario, every reactant molecule would completely convert into product molecules, resulting in the theoretical yield. However, this is rarely observed in reality. Several reasons contribute to the discrepancy between actual and theoretical yields.

Firstly, reactions may not proceed to completion. This occurs when not all reactant molecules are converted into products. Factors such as the reaction reaching equilibrium before all reactants are consumed, or the presence of reactants that are inaccessible for reaction—such as those that are inadequately mixed or present in large, solid clumps—can hinder the reaction.

Secondly, side reactions can occur. These are unintended reactions that take place alongside the primary reaction, consuming some reactants to form alternative products. Side reactions are particularly likely when the reactants or products are reactive substances that may interact with each other or with external elements such as air or moisture.

Lastly, product loss during purification processes can contribute to a lower actual yield. For instance, if the product is a solid that is filtered from a liquid, some of it may inadvertently remain in the liquid. In cases where the product is a gas, a portion may escape into the atmosphere. Even with liquids, some product may be left behind in the container when it is transferred.

In summary, while the theoretical yield is determined by the stoichiometry of the balanced chemical equation, the actual yield is frequently lower due to practical challenges such as incomplete reactions, side reactions, and product loss during purification.