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
Less waste, lower raw material cost, more sustainable process
Products re-form reactants — reaction doesn't go to completion
Actual yield 38 g, theoretical yield 50 g — (38÷50) × 100
Product sticks to glassware or lost during filtration and transfer
⚽ 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.
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?
⭐ How Well Do You Understand This Topic?
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