1.5.5Principles of Chemistry
Calculate Percentage Yield — Edexcel International GCSE Science Double Award 4SD0 - Chemistry Revision Notes & Topical Questions
5 min read 12 parts Lesson 26 of 50
Part 1 of 5Two kinds of yield
Yield
- Yield The yield of a reaction is the mass of product that is obtained from it.
- Industrial chemists want a high yield because wasted product means wasted money, wasted raw materials and more waste to dispose of.
Theoretical yield
- Theoretical yield This is the maximum mass of product that could form if every particle of the limiting reactant reacted exactly as the balanced equation says.
- It is a calculated value, not a measured one, and it assumes perfect conditions with nothing lost.
Actual yield
- Actual yield This is the mass of product you really collect, weighed at the end of the experiment or process.
- It is nearly always smaller than the theoretical yield, and it can never legitimately be larger.
Part 2 of 5Why the actual yield is less than 100%
Reasons for a lower yield
Reversible reactions
- Reversible reactions In a reversible reaction the products keep turning back into reactants, so the reaction never goes fully to completion.
Losses during handling
- Losses during handling Some product is left on filter paper, stuck to glassware, or spilt when it is moved from one container to another.
Losses during separation
- Losses during separation Filtration, crystallisation, distillation and purification all lose a little of the product.
Side reactions
- Side reactions Reactants may react with oxygen in the air or with impurities, forming unwanted by-products instead of the desired substance.
Incomplete reaction
- Incomplete reaction Some reactant may be left unreacted, stuck in the apparatus, if the reaction is too slow or is stopped too early.
Part 3 of 5The percentage yield equation
| Quantity | Equation | Notes |
|---|---|---|
| Percentage yield | percentage yield = (actual yield ÷ theoretical yield) × 100 | Both masses must be in the same unit, such as grams. |
| Moles | n = m ÷ Mr | Used to convert the starting mass into moles of reactant. |
| Mass | m = n × Mr | Used to convert moles of product back into a theoretical mass. |
| Result | 0% to 100% | A value above 100% means an error, often a product that is still wet or impure. |
Part 4 of 5Worked example: from mass to percentage yield
1
Question: 4.0 g of magnesium burns in excess oxygen: 2Mg + O2 → 2MgO. A student collects 5.8 g of magnesium oxide. Find the percentage yield. (Ar: Mg = 24, O = 16.)
2
Moles of magnesium = 4.0 ÷ 24 = 0.167 mol.
3
The equation shows 2 mol Mg gives 2 mol MgO, a 1 : 1 ratio, so the theoretical amount of MgO is also 0.167 mol.
4
Mr of MgO = 24 + 16 = 40, so theoretical yield = 0.167 × 40 = 6.67 g.
5
Percentage yield = (5.8 ÷ 6.67) × 100 = 87%.
6
Conclusion: 87% of the maximum possible magnesium oxide was collected; the rest was lost, for example as smoke escaping when the crucible lid was lifted.
Part 5 of 5Why percentage yield matters
Measuring success
- Measuring success Percentage yield is a simple way of judging how efficient a chemical process is.
Choosing a reaction pathway
- Choosing a reaction pathway A product can often be made by several routes, each a sequence of reactions called a reaction pathway.
- Companies compare the overall percentage yields of different pathways, together with cost, speed and hazards, to pick the best one.
Economics and waste
- Economics and waste A higher yield means lower costs, higher profit and less waste, which is why industry works hard to raise it.
Key terms
| Term | Meaning |
|---|---|
| Yield | The mass of product obtained from a chemical reaction. |
| Theoretical yield | The maximum mass of product calculated from the balanced equation, assuming complete reaction and no losses. |
| Actual yield | The mass of product that is really obtained and measured. |
| Percentage yield | The actual yield divided by the theoretical yield, multiplied by 100. |
| Reaction pathway | A sequence of reactions that together turn starting materials into the required product. |
| Side reaction | An unwanted reaction that uses up reactants to make by-products. |
Avoid these mistakes
1
Putting the theoretical yield on top of the fraction. The actual yield is always the numerator, so the answer cannot exceed 100%.
2
Forgetting to use the balanced equation ratio when finding moles of product, especially when the ratio is not 1 : 1.
3
Saying the yield is below 100% because the 'law of conservation of mass is broken'. Mass is conserved; the product is simply lost or never forms.
4
Mixing units, for example using kilograms for one yield and grams for the other, which gives a wrong percentage.
5
Using the mass of an excess reactant to calculate the theoretical yield. It must be based on the limiting reactant.
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