AP Chemistry Limiting Reactant & Percent Yield
How to find the limiting reactant, then calculate theoretical and percent yield, with full worked examples.
Limiting reactant and percent yield sit inside Topic 4.5 (Stoichiometry) of Unit 4, and they almost always show up together: you find the limiting reactant first, use it to calculate a theoretical yield, then compare that to an actual yield to get a percent yield. Here's the full method, worked step by step, plus the AP scoring detail most guides skip entirely.
Step 1: Find the Limiting Reactant
Whenever a problem gives you starting amounts of two reactants, don't assume either one is fully used up, you have to check. The method: convert both given amounts to moles, then divide each by its coefficient in the balanced equation. Whichever reactant gives the smaller result is the limiting reactant.
Worked example: N₂ + 3 H₂ → 2 NH₃. You start with 28.0 g N₂ and 9.0 g H₂. Which one runs out first?
- 28.0 g N₂ ÷ 28.0 g/mol = 1.00 mol N₂; divided by its coefficient (1) = 1.00
- 9.0 g H₂ ÷ 2.0 g/mol = 4.5 mol H₂; divided by its coefficient (3) = 1.5
1.00 is smaller than 1.5, so N₂ is the limiting reactant and H₂ is in excess. (If the numbers on both sides come out equal instead, like 28.0 g N₂ against exactly 6.0 g H₂, the reactants are in a perfect stoichiometric ratio and neither one is technically in excess.)
Step 2: Calculate Theoretical Yield
Once you know the limiting reactant, theoretical yield comes from a single dimensional-analysis chain: moles of limiting reactant → moles of product (using the mole ratio from the balanced equation) → grams of product (using the product's molar mass).
Continuing the example (1.00 mol N₂ limiting): 1.00 mol N₂ × (2 mol NH₃ / 1 mol N₂) = 2.00 mol NH₃. Convert to grams: 2.00 mol × 17.0 g/mol = 34.0 g NH₃ theoretical yield.
Step 3: Calculate Percent Yield
Percent yield compares what you actually collected to that theoretical maximum: percent yield = (actual yield ÷ theoretical yield) × 100.
If the reaction above actually produced 29.8 g NH₃ in the lab: 29.8 g ÷ 34.0 g × 100 = 87.6% yield. That's a realistic number, real reactions lose product to side reactions, reactions that don't go fully to completion, and ordinary physical losses during transfer and purification. A calculated yield over 100% almost always means a measurement or rounding error somewhere, not a reaction that beat its own theoretical ceiling.
Don't Forget: Leftover Excess Reactant
A second question type asks how much of the excess reactant is left over, not just how much product forms. Use the limiting reactant to calculate how much of the excess reactant actually got used, then subtract that from the starting amount.
Example: 1.00 mol N₂ (limiting) reacts with 4.5 mol H₂ available. N₂ consumes 1.00 mol × (3 mol H₂ / 1 mol N₂) = 3.00 mol H₂. Leftover H₂ = 4.5 mol − 3.00 mol = 1.5 mol H₂ remaining, unreacted.
Common Mistakes
- Picking the limiting reactant based on which starting mass is smaller, instead of converting to moles and dividing by the coefficient first. The smaller mass isn't always the limiting reactant.
- Using the excess reactant's amount to calculate theoretical yield instead of the limiting reactant's, which always overstates the yield.
- Forgetting to convert theoretical and actual yield to the same units before dividing for percent yield.
- Reporting a percent yield over 100% without questioning it, rather than treating it as a sign to recheck the calculation.
Related Resources
- Unit 4 Review: Chemical Reactions
- AP Chemistry Significant Figures Rules
- AP Chemistry Reference Sheet
- AP Chemistry FRQs Explained
- AP Chemistry FRQ Calculator
Frequently Asked Questions
What is a limiting reactant?
The limiting reactant is whichever reactant runs out first, and it's the one that determines how much product a reaction can actually form. Every other reactant is in excess: some of it is left over once the reaction stops, because there's nothing left for it to react with.
What is theoretical yield?
Theoretical yield is the maximum amount of product a reaction could form, calculated entirely from the limiting reactant using the mole ratios in the balanced equation. It's a calculated ceiling, not a measured result, real reactions almost never reach it.
What is the percent yield formula?
Percent yield = (actual yield / theoretical yield) x 100. Actual yield is what you actually collect in the lab (given in the problem or measured); theoretical yield is the calculated maximum from the limiting reactant. Both values need the same units before you divide.
Why is percent yield always less than 100%?
Real reactions lose product to side reactions, reactions that don't go to completion, and physical losses during transfer, filtration, or purification. A percent yield over 100% almost always signals a measurement or calculation error, not a genuinely better-than-perfect reaction, most often leftover impure product that wasn't fully dried or purified.
Does AP Chemistry give partial credit if I use the wrong limiting reactant?
Often yes. AP Chemistry FRQ rubrics commonly award consistent-with credit on multi-step calculations: if you identify the wrong limiting reactant but then correctly carry that value through the rest of the math, you typically lose the point for the identification step but can still earn the calculation points, since the rest of your work is internally consistent. It's not automatic and depends on the specific rubric, but showing your full method is what makes that credit possible at all.
This page is not affiliated with or endorsed by College Board. AP® is a trademark registered by the College Board.