5.12C - Why actual yield is usually lower than theoretical yield

Actual and theoretical yield

In chemistry, a yield is the amount of desired product made in a reaction. The key comparison is between the amount that could be made in theory and the amount actually collected from the experiment.

This might be a mass in g for a solid product, or another measured amount if the product is collected in a different form. It should refer to the product, not to the reactants.

Theoretical yield is not just a guess. It is the calculated maximum, assuming the reaction goes as far as possible to form the desired product and the product is collected without loss.

For example, a calculation might show that a reaction could make 5.00 g of a dry salt. If the student actually collects 4.10 g, the theoretical yield is 5.00 g and the actual yield is 4.10 g. The actual yield is lower because real experiments do not usually match the ideal calculation.

Incomplete reactions

One reason actual yield is lower is that the reaction may be incomplete.

If some of the reactant that should have reacted remains, less desired product has formed than the calculation assumed. A deliberately excess reactant can remain even when the reaction is complete: leftover excess alone does not prove incomplete reaction. A reversible reaction may also reach equilibrium before all the available reactant forms product. The theoretical yield is based on the reaction producing as much desired product as possible, but the actual yield only counts what is really obtained.

The exam answer needs the link: some reactant is not converted, so less desired product forms. Writing only "the reaction was incomplete" is often too thin unless the effect on product is clear.

Side reactions

A third reason actual yield is lower is that competing, unwanted reactions may happen.

In a side reaction, some reactant is used to make an unwanted product instead of the desired product. That reduces the amount of desired product available to collect. The total amount of material in the reaction mixture may still be large, but the actual yield of the desired product is lower.

For a simple way to think about it:

reactants -> desired product
reactants -> unwanted product

The second route competes with the first route. A good answer identifies a competing reaction (a named reaction is not required) and says that reactants are used up making unwanted products, so less desired product is made.

Practical losses

Another reason actual yield is lower is that some product is made but is not collected.

Practical losses happen during the experimental work after or during the reaction. For a solid product, some crystals might stay on filter paper, stick to the side of a beaker, be left behind during transfer, or be lost during washing and drying. The product existed, but it did not end up in the final measured sample.