Actual yield = mass of product actually obtained (g)
Theoretical yield = maximum possible mass calculated from the equation (g)
EXAMPLE:
Theoretical yield of calcium carbonate = 50 g
Actual yield collected = 38 g
Percentage yield = (38 ÷ 50) × 100 = 76%
A percentage yield of 100% would mean no product was lost — impossible in practice.
A high percentage yield is desirable — means less waste and lower costs.
INDUSTRIAL IMPORTANCE:
Chemical manufacturers aim to maximise yield to:
Reduce waste and environmental impact.
Reduce costs — fewer raw materials wasted.
Improve profitability of the process.
Why Maximising Yield Matters
ECONOMIC REASONS:
A low yield means raw materials are wasted — higher production costs.
Manufacturers need high yields to make processes economically viable.
ENVIRONMENTAL REASONS:
Wasted reactants may be polluting or energy-intensive to produce.
High yield means less waste entering the environment.
SUSTAINABILITY:
Raw materials are often finite resources — minimising waste conserves them.
EXAMPLE — Haber process:
The reaction N₂ + 3H₂ ⇌ 2NH₃ is reversible.
At equilibrium, only about 15% conversion to ammonia per pass.
Unreacted nitrogen and hydrogen are recycled — overall yield improved to ~98%.
PRACTICAL TIP — improving yield:
Ensure complete reactions: sufficient time, temperature, catalyst.
Minimise product loss during separation and purification.
Recycle unreacted starting materials where possible.
⚠️ Common Mistake
The theoretical yield must be calculated from the balanced equation — not just assumed. Use moles: find moles of limiting reactant → use molar ratio → calculate mass of product. Then divide actual mass by this calculated theoretical mass.
% yield = (actual ÷ theoretical) × 100. Always less than 100% in practice. Reasons: reversible reactions, side reactions, practical losses, impurities. High yield = less waste, lower cost, more sustainable. Chemistry-only spec point.
🎯 Matching Activity — Percentage Yield
Match each scenario to the correct percentage yield or reason for low yield. — drag the symbols on the right to match the component names on the left.
76%
Drop here
Reversible reaction
Drop here
Practical losses
Drop here
High yield is desirable
Drop here
Product sticks to glassware or lost during filtration and transfer
Actual yield 38 g, theoretical yield 50 g — (38÷50) × 100
Less waste, lower raw material cost, more sustainable process
Products re-form reactants — reaction doesn't go to completion
⚽ FIFA Worked Examples
Percentage Yield
A student expects to make 8.0 g of copper sulfate but only collects 6.2 g. Calculate the percentage yield.
Calculate theoretical yield using moles: convert reactant mass → moles → use molar ratio from equation → convert to product mass. Calculate percentage yield from these values. Evaluate the economic and environmental importance of maximising percentage yield. Identify the limiting reactant in a reaction and use it to calculate maximum theoretical yield.
🔬 Triple Science Only
Percentage yield (4.3.3.1) is chemistry-only — not in Combined Science. Students must calculate percentage yield and explain why actual yield is less than theoretical yield.
🎯 Test Yourself
Question 1 of 2
1. A reaction has a theoretical yield of 20 g but only 14 g is obtained. What is the percentage yield?
2. Why is the actual yield of a reaction almost always less than the theoretical yield?
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