Understanding Limiting Reagents
What is a Limiting Reagent?
The limiting reagent (or limiting reactant) is the reactant that is completely consumed first in a chemical reaction. It determines the maximum amount of product that can be formed.
Key Concept:
The reactant that produces the LEAST amount of product is the limiting reagent.
Step-by-Step Process
Step 1: Write the balanced chemical equation
Example: 2H₂ + O₂ → 2H₂O
Step 2: Convert all given amounts to moles
• If given grams: moles = grams ÷ molar mass
• If given moles: use directly
Step 3: Calculate mole ratio for each reactant
• Divide moles available by coefficient in balanced equation
• The reactant with the SMALLEST ratio is the limiting reagent
Step 4: Calculate theoretical yield
• Use moles of limiting reagent
• Apply stoichiometry from balanced equation
• Convert to desired unit (g, mol, etc.)
Step 5: Calculate excess reagent remaining
• Determine how much was used based on limiting reagent
• Subtract from initial amount
Example Problem
Problem: 4.0 g of H₂ reacts with 32.0 g of O₂. Find the limiting reagent and theoretical yield of H₂O.
Balanced equation: 2H₂ + O₂ → 2H₂O
Step 1 - Convert to moles:
• H₂: 4.0 g ÷ 2.016 g/mol = 1.98 mol
• O₂: 32.0 g ÷ 32.00 g/mol = 1.00 mol
Step 2 - Calculate mole ratios:
• H₂: 1.98 mol ÷ 2 = 0.99
• O₂: 1.00 mol ÷ 1 = 1.00
Result: H₂ is the limiting reagent (smallest ratio)
Step 3 - Theoretical yield:
• From equation: 2 mol H₂ → 2 mol H₂O
• So: 1.98 mol H₂ → 1.98 mol H₂O
• Mass: 1.98 mol × 18.015 g/mol = 35.7 g H₂O
Percent Yield
Percent yield compares actual yield (what you got in the lab) to theoretical yield (maximum possible):
Percent Yield Formula:
% Yield = (Actual Yield ÷ Theoretical Yield) × 100%
Why is percent yield less than 100%?
- Incomplete reactions
- Side reactions producing unwanted products
- Product lost during purification/transfer
- Measurement errors
Common Mistakes to Avoid
- ❌ Forgetting to balance the equation first
- ❌ Using grams instead of moles for comparison
- ❌ Not dividing by stoichiometric coefficients
- ❌ Assuming the reactant with less mass is limiting
- ✅ Always convert to moles and use mole ratios!
Real-World Applications
- Industrial Chemistry: Minimize waste by using excess of cheaper reactant
- Pharmaceutical Manufacturing: Calculate exact amounts needed for drug synthesis
- Environmental Chemistry: Predict pollutant formation in combustion
- Food Chemistry: Optimize ingredient ratios in recipes