GCSE Chemistry revision

Percentage yield

Yield and atom economy of chemical reactions (chemistry only)

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The theoretical yield is the amount of product you would expect to get based on calculations, assuming all reactants are converted.
The actual yield is the amount of product you actually obtain when you carry out the reaction in practice.

When 2 g of hydrogen reacts with 16 g of oxygen, the theoretical yield of water is 18 g, since 2 + 16 = 18.
In practice, the actual yield is often lower than the theoretical yield — for example, only 15 g of water might be produced.

Some reactants may remain unreacted at the end of a reaction, reducing the amount of product formed.
A slow reaction may not have had sufficient time to go to completion, leaving unreacted starting materials.
Reversible reactions reach a state of equilibrium and never go to completion — for example, the reaction of nitrogen and hydrogen to form ammonia, where some ammonia continuously breaks back down into nitrogen and hydrogen.

3. Reason 1: Reactants Not Fully Reacting AQA

Some reactants may remain unreacted at the end of a reaction, reducing the amount of product formed.
A slow reaction may not have had sufficient time to go to completion, leaving unreacted starting materials.
Reversible reactions reach a state of equilibrium and never go to completion — for example, the reaction of nitrogen and hydrogen to form ammonia, where some ammonia continuously breaks back down into nitrogen and hydrogen.

Side reactions occur when reactants react to form an unintended product instead of the desired one, reducing the yield of the target product.
For example, in the production of ammonia, some nitrogen may react with oxygen in the air to produce nitrogen dioxide instead, meaning less ammonia is formed than expected.

4. Reason 2: Side Reactions AQA

Side reactions occur when reactants react to form an unintended product instead of the desired one, reducing the yield of the target product.
For example, in the production of ammonia, some nitrogen may react with oxygen in the air to produce nitrogen dioxide instead, meaning less ammonia is formed than expected.

Product can be lost during practical procedures, such as gaseous products escaping into the atmosphere.
During filtration, liquid product can be lost as some remains in the original test tube, on the solid residue, or absorbed by the filter paper.
If the solid is the desired product, some may be left behind on the filter paper when attempting to scrape it off, reducing the actual yield.

5. Reason 3: Loss of Product During Processing AQA

Product can be lost during practical procedures, such as gaseous products escaping into the atmosphere.
During filtration, liquid product can be lost as some remains in the original test tube, on the solid residue, or absorbed by the filter paper.
If the solid is the desired product, some may be left behind on the filter paper when attempting to scrape it off, reducing the actual yield.

Percentage yield tells us what proportion of the theoretical yield was actually obtained, expressed as a percentage.
The formula for percentage yield is: Percentage yield = actual yield / theoretical yield × 100 Percentage yield ranges from 0% (no product obtained) to 100% (all predicted product obtained).
Using the earlier example, a theoretical yield of 18 g and an actual yield of 15 g gives a percentage yield of 15 / 18 × 100 = 83.3%.

6. Calculating Percentage Yield AQA

Percentage yield tells us what proportion of the theoretical yield was actually obtained, expressed as a percentage.
The formula for percentage yield is: Percentage yield = actual yield / theoretical yield × 100 Percentage yield ranges from 0% (no product obtained) to 100% (all predicted product obtained).
Using the earlier example, a theoretical yield of 18 g and an actual yield of 15 g gives a percentage yield of 15 / 18 × 100 = 83.3%.

Yield is the mass of product formed in a reaction, and can be expressed as actual yield, theoretical yield, or percentage yield.
A percentage yield of 100% is ideal but rarely achieved in practice due to factors such as incomplete reactions, side reactions, and product loss.