Problem 107 - Chemistry The Science in Context

Question

Give two reasons why the actual yield from a chemical reaction is usually less than the theoretical yield.

Step-by-Step Solution

Answer

Question: Explain two reasons why the actual yield from a chemical reaction is usually less than the theoretical yield. Answer: One reason is that chemical reactions might not go to completion due to equilibrium limitations or side reactions, leading to a percentage of reactants remaining unreacted. Another reason is the loss of product during purification or isolation steps, where a portion of the desired product might be lost or degraded.

Step 1: Reason 1: Incomplete Reactions

One reason for the discrepancy between actual and theoretical yield is the fact that chemical reactions might not go to completion. In some cases, the reactants do not fully convert into the desired products due to equilibrium limitations or the presence of side reactions. In such cases, a percentage of the reactants remains unreacted, and as a result, the actual yield is lesser than the calculated theoretical yield.

Step 2: Reason 2: Loss of Product during Purification

Another reason for the lower actual yield is the loss of product during purification or isolation steps. Most chemical reactions produce a mixture of products that need to be separated and purified. During these processes, a portion of the desired product might be lost or degraded, leading to a decrease in the actual yield compared to the theoretical yield.

Key Concepts

Incomplete Chemical Reactions

When performing a chemical reaction, students often expect to obtain the full amount of product as predicted by the stoichiometry of the reaction, called the theoretical yield. However, in practice, this yield is rarely achieved due to incomplete chemical reactions. This occurs when not all reactant molecules react to form the intended product. Equilibrium limitations can prevent the complete conversion of reactants when the reaction is reversible. In these cases, the reaction achieves a balance between reactants and products where both are present at specific concentrations when no further changes occur.

For example, if a synthesis reaction is supposed to produce \( 100 \) grams of product, but only \( 75 \) grams are obtained, the reaction may not have gone to completion. Even if no errors were made during the experiment, aspects like reaction time, temperature, pressure, and the presence of catalysts can alter the extent of the reaction. Therefore, understanding that chemical reactions often do not go to completion is crucial for students who are comparing actual and theoretical yields.

Product Purification Losses

After a chemical reaction occurs, the resulting mixture often contains by-products alongside the target compound. To isolate the desired product, purification steps like filtration, distillation, or chromatography are required. However, during these purification processes, some of the desired product may be inadvertently lost. This can occur due to several factors:

• Product sticking to the filter material or the walls of the container.
• Partial decomposition of the product during purification techniques that involve heat (like distillation).
• Inefficient separation methods leading to the discarding of some product along with by-products or solvents.

These losses can be minimized but not always entirely prevented, contributing to a lower actual yield. This reduction can be particularly significant if the product is volatile, sensitive to heat, or prone to degradation. Whether using mechanical, thermal, or chemical methods for purification, it's important for students to understand that some product will likely be lost at this stage, highlighting a difference between actual and theoretical yields.

Chemical Reaction Equilibrium

Many chemical reactions are reversible, meaning they can proceed in both the forward and reverse directions. When the rate of the forward reaction equals the rate of the reverse reaction, the system reaches a state known as chemical equilibrium. At this point, the relative concentrations of reactants and products remain constant. The position of equilibrium depends on various factors like temperature, pressure, and concentration, according to Le Chatelier's Principle.

For example, if a reaction at equilibrium favors the reactants, it means that even though some product is formed, a significant amount of reactants remains unconverted. This affects the actual yield because the maximum theoretical yield assumes complete conversion of reactants, which doesn't happen at equilibrium. Consequently, when predicting the amount of product formed, considering the reaction's equilibrium position is vital for a realistic estimate of the actual yield.

Side Reactions

Side reactions are unintended reactions that occur simultaneously with the desired reaction. They can consume reactants or produce additional products that are not part of the intended reaction pathway. The presence of impurities, incorrect reactant ratios, or uncontrolled reaction conditions can lead to these competing reactions.

For instance, suppose a reaction is designed to produce only one product from specific reactants. If a secondary reaction starts producing an unwanted compound, the reactants are divided between forming the desired and undesired products. This division reduces the amount of desired product, thus decreasing the actual yield. Identifying and minimizing side reactions through careful control of reaction conditions and purification of reactants is essential in optimizing yields and achieving results closer to the theoretical prediction.