5.7 - Percentage Yield

Why a reaction gives less product than expected

• The mass of product a reaction should make can be calculated from the balanced equation and the masses of the reactants used.
• The mass actually obtained is almost always less than this theoretical amount.

1. The reaction doesn't finish

• A reaction stopped before it's complete leaves unreacted starting material.
• A reversible reaction can't fully complete because products turn back into reactants (e.g. the Haber process making ammonia).

2. Side reactions occur

Reactants can react with gases in the air or with impurities, forming unintended products instead of the wanted one.

3. Product is lost during separation and purification

• Solid product is left behind stuck to filter paper or container walls.
• Liquid product is lost wetting the filter paper or remaining on glassware.
• A gas product can escape into the air before it's collected.

Theoretical yield, actual yield and percentage yield

• Theoretical yield - The maximum mass of product that could be formed from a given amount of reactant, calculated from the balanced equation.
• Actual yield - The mass of product actually obtained and weighed when the reaction is carried out.
• Percentage yield - The actual yield expressed as a percentage of the theoretical yield.

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Formula for percentage yield

Both yields must be in the same unit (usually grams) before calculating.

What percentage yield tells you

• 100% - All theoretical product was obtained.
• 0% - No reactants were converted to product.
• A yield above 100% is impossible by conservation of mass, and usually means the product is still wet with solvent or an arithmetic error was made.
• A low percentage yield means reactants were wasted, which increases costs and reduces sustainability in industrial processes.

Worked example - Calculating a percentage yield

A reaction to make copper(II) sulfate has a theoretical yield of 18.0 g. The mass of copper(II) sulfate obtained is 15.0 g. Calculate the percentage yield, to three significant figures.