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S1.1 - Introduction to the particulate nature of matter S1.2 - The nuclear atom S1.3 - Electron configurations S1.4 - Counting particles by mass - The mole S1.5 - Ideal gases S2.1 - The ionic model S2.2 - The covalent model S2.3 - The metallic model S2.4 - From models to materials S3.1 - The periodic table - Classification of elements S3.2 - Functional groups - Classification of organic compounds R1.1 - Measuring enthalpy changes R1.2 - Energy cycles in reactions R1.3 - Energy from fuels R1.4 - Entropy and spontaneity AHL R2.1 - How much? The amount of chemical change R2.2 - How fast? The rate of chemical change R2.3 - How far? The extent of chemical change R3.1 - Proton transfer reactions R3.2 - Electron transfer reactions R3.3 - Electron sharing reactions R3.4 - Electron-pair sharing reactions

R2.1 - How much? The amount of chemical change

2.1.1 Chemical Equations and State Symbols 2.1.2 Stoichiometry and Mole Reaction 2.1.3 + 2.1.4 Limiting Reactants and Yield 2.1.5 Atom Economy

Limiting Reactants and Percentage Yield

Specification Reference R2.1.3 & R2.1.4

Quick Notes

  • The limiting reactant is the reactant that is completely used up and determines the theoretical yield in a reaction.
  • The excess reactant is the one that remains after the reaction is complete.
  • Theoretical yield is the calculated maximum amount of product based on the limiting reactant.
  • Experimental yield is the amount actually obtained in the lab.
  • Percentage yield = (experimental yield / theoretical yield) × 100

Full Notes

What is the Limiting Reactant?

In a chemical reaction, the limiting reactant is the one that runs out first. It limits how much product can form.

The excess reactant is not fully used up and some will be left over.

Identifying the Limiting Reactant

Step-by-step method:

Worked Example

If 28.0 g N2 and 6.0 g H2 are mixed together and react according to the equation N2 + 3H2 → 2NH3, which is the limiting reagent?

  1. Molar mass of N2 = 28.02 g mol⁻¹ → n = 28.0 / 28.02 = 1.00 mol
  2. Molar mass of H2 = 2.02 g mol⁻¹ → n = 6.0 / 2.02 ≈ 2.97 mol
  3. From the equation, 1 mol N2 reacts with 3 mol H2.
  4. We have just enough N2 for 3.00 mol H2, but we only have 2.97 mol H2.
  5. Therefore, H2 is the limiting reactant.

Theoretical vs Experimental Yield

A theoretical yield is the maximum amount of product that could be obtained in a reaction, based on the amount of the limiting reactant.

An experimental (or actual) yield is the amount of product that is actually obtained in the reaction.

Theoretical and experimental yields are often compared to each other using percentage yield.

IB Chemistry percentage yield formula: percentage yield = (actual yield / theoretical yield) × 100.
Worked Example

In the reaction Mg + 2HCl → MgCl2 + H2, if 2.40 g of Mg is reacted with excess HCl, and only 5.80 g of MgCl2 is obtained, what is the percentage yield?

IB Chemistry worked example for Mg reacting with HCl showing theoretical yield of MgCl2, actual yield 5.80 g, percentage yield = 61.7%.

Why Percentage Yield is Always Less Than 100%

In reality, percentage yields are rarely 100% because of:

Tool 1, Inquiry 1, 2, 3 – Linked Course Question

What errors may cause the experimental yield to be
i) higher and ii) lower than the theoretical yield?

i) Higher than theoretical yield:

  • Contamination of the product (e.g., impurities or residual solvent)
  • Incomplete drying of a solid product (extra water adds mass)
  • Instrumental or weighing errors (e.g., balance not zeroed)

ii) Lower than theoretical yield:

  • Incomplete reactions or side reactions
  • Loss of product during transfer, filtration, or purification
  • Evaporation of volatile substances
  • Impurities in reactants reducing efficiency

Summary