Equilibrium Constant: Kp | A-level Chemistry | OCR, AQA, Edexcel

Introduction

This tutorial provides a comprehensive guide to understanding the equilibrium constant (Kp) in chemistry, specifically focusing on partial pressures and mole fractions. This concept is essential for A-level Chemistry students studying equilibrium reactions involving gases. By the end of this tutorial, you will be able to calculate partial pressures, determine mole fractions, and understand how to apply the equilibrium constant in various situations.

Step 1: Understanding Partial Pressures

Partial pressure is the pressure that a specific gas would exert if it occupied the entire volume alone. Here’s how to understand and calculate partial pressures:

• Total Pressure: In a mixture of gases, the total pressure is the sum of the partial pressures of each gas.

• Symbol: We often use p(A) to denote the partial pressure of gas A.

• Calculation: The partial pressure can be calculated using the formula:

  p(A) = Xa x Total Pressure
  

Step 2: Calculating Mole Fractions

Mole fractions help quantify the amount of a specific gas in a mixture. Here’s how to calculate it:

• Formula: The mole fraction (Xa) is calculated using:

  Xa = number of moles of substance A / total number of moles of all substances
  

• Example Calculation:

  • Given: 0.14 moles of O2, 0.52 moles of N2, and 0.01 moles of other gases.

  • Total moles = 0.14 + 0.52 + 0.01 = 0.67 moles.

  • Mole fraction of oxygen:

    Xa = 0.14 / 0.67 ≈ 0.209
    

Step 3: Calculating Partial Pressures from Mole Fractions

To find the partial pressure of a gas in a mixture, follow these steps:

1. Calculate the mole fraction as outlined in Step 2.
2. Multiply the mole fraction by the total pressure of the gas mixture.

• Example Calculation:

  • Given total pressure = 140 kPa.

  • Using the mole fraction of oxygen calculated previously (≈ 0.209):

    p(O2) = 0.209 x 140 kPa ≈ 29.26 kPa
    

Step 4: Understanding the Equilibrium Constant Kp

For equilibrium reactions involving gases, we can express the equilibrium constant (Kp) in terms of partial pressures.

• General Form: For a reaction:

  aA(g) + bB(g) ⇌ cC(g) + dD(g)
  

  The equilibrium constant Kp is given by:

  Kp = (pC^c * pD^d) / (pA^a * pB^b)
  

• Example Calculation:

  • Given partial pressures: p(N2O4) = 50 kPa and p(NO2) = 68 kPa.

  • For the reaction:

    N2O4(g) ⇌ 2NO2(g)
    

    The Kp would be calculated as:

    Kp = (p(NO2)^2) / (p(N2O4)) = (68^2) / 50
    

Step 5: Applying Kp in Heterogeneous Reactions

When dealing with heterogeneous reactions (involving different states of matter), remember the following:

• Exclusions in Kp:
  • Do not include pure solids or pure liquids in the Kp expression.
  • Only gases are included in the Kp calculations.

Conclusion

Understanding Kp and its relation to partial pressures and mole fractions is crucial for mastering equilibrium concepts in chemistry. Remember to:

• Calculate mole fractions accurately.
• Use the correct formulas for partial pressures and Kp.
• Exclude solids and liquids when calculating Kp in heterogeneous reactions.

With these skills, you can confidently tackle equilibrium problems in your A-level Chemistry studies. For further practice, try working through sample problems or past exam questions related to Kp.