Percentage Yield | A-level Chemistry | OCR, AQA, Edexcel

Introduction

This tutorial explains the concept of percentage yield in chemistry, its importance, and how to identify limiting reagents in a chemical reaction. Understanding these principles is crucial for evaluating the efficiency of synthetic procedures in A-level Chemistry.

Step 1: Understand Percentage Yield

Percentage yield measures the effectiveness of a chemical reaction by comparing the actual yield to the theoretical yield.

• Formula: [ \text{Percentage Yield} = \left( \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \right) \times 100% ]

• Key Points:

  • Actual yield refers to the amount of product obtained from a reaction.
  • Theoretical yield is the maximum amount of product expected based on stoichiometric calculations.
  • A percentage yield of 100% indicates that all reactants were converted into products, which is rare in practice.

Step 2: Recognize the Importance of Percentage Yield

Understanding percentage yield helps assess the success of a chemical reaction.

• Importance:
  • Indicates how much of the reactants have been converted into products.
  • Percentage yield helps identify inefficiencies in a reaction, such as:
    • Formation of unwanted by-products from side reactions.
    • Unreacted reactants remaining in the vessel.
    • Loss of product during recovery or transfer.
    • Reactions not proceeding to completion due to equilibrium.

Step 3: Identify Limiting Reagents

In many reactions, one reactant will limit the amount of product that can be formed. This is known as the limiting reagent.

• Definition:

  • The limiting reagent is the reactant that is completely consumed first, thus determining the amount of product formed.

• Example: If you have a recipe that requires:

  • 3 eggs
  • 100g flour And your cupboard contains:
  • 10 eggs
  • 1kg flour The limiting reagent is eggs because they will run out before the flour.

Step 4: Solve a Worked Example

To apply the concepts learned, consider the reaction of nitrogen and hydrogen to form ammonia.

• Reaction Equation: [ N_2 + 3H_2 \rightarrow 2NH_3 ]

• Given:

  • 0.5g of Nitrogen
  • 0.6g of Hydrogen

• Determine the Limiting Reagent:

  1. Calculate the molar masses:
     • Nitrogen (N2): 28 g/mol
     • Hydrogen (H2): 2 g/mol
  2. Convert grams to moles:
     • Moles of Nitrogen = 0.5 g / 28 g/mol = 0.0179 moles
     • Moles of Hydrogen = 0.6 g / 2 g/mol = 0.3 moles
  3. Use the stoichiometry of the reaction:
     • The reaction requires 3 moles of H2 for every mole of N2.
     • For 0.0179 moles of N2, you would need: [ 0.0179 \times 3 = 0.0537 \text{ moles of H2} ]
     • Since 0.3 moles of H2 is available, nitrogen is the limiting reagent.

Conclusion

In summary, understanding percentage yield and limiting reagents is essential for analyzing chemical reactions. Remember to calculate both actual and theoretical yields to determine percentage yield and identify limiting reagents to predict the amount of product formed. Next, practice these concepts with different chemical equations to reinforce your understanding.