Why might the actual yield be less than the theoretical yield?

The actual yield might be less than the theoretical yield due to incomplete reactions, side reactions, or loss of product during recovery.

In a chemical reaction, the theoretical yield is the maximum amount of product that can be produced from a given amount of reactants. However, in reality, the actual yield is often less than the theoretical yield. This discrepancy can be attributed to several factors.

Firstly, the reaction may not go to completion. This means that not all of the reactants are converted into the desired product. This could be due to the reaction reaching equilibrium before all the reactants have reacted, or because some of the reactants are used up in side reactions.

Secondly, side reactions may occur. These are unintended reactions that compete with the main reaction for the reactants. Side reactions often produce unwanted products, reducing the amount of reactants available for the main reaction and hence decreasing the actual yield.

Lastly, there can be loss of product during the recovery process. After the reaction has taken place, the desired product often needs to be separated from the reaction mixture. This process can be complex and may result in some of the product being lost, for example, due to evaporation or adherence to the walls of the container.

In addition, factors such as measurement errors, impure reactants, or incorrect stoichiometric calculations can also lead to a lower actual yield. Therefore, while the theoretical yield provides an ideal outcome based on stoichiometric calculations, the actual yield reflects the realities of performing chemical reactions in a laboratory setting.