Physical Science Matric Revision: Chemical Equilibrium

Chemical Equilibrium Revision Notes

Introduction

Chemical equilibrium is a fundamental concept in chemistry, particularly in understanding how reactions behave in closed systems. This concept is crucial for predicting how changes in conditions like concentration, temperature, and pressure affect the composition of reactants and products.

Learning Objectives:

  • Understand what chemical equilibrium is.
  • Learn to apply Le Chatelier’s Principle.
  • Calculate the equilibrium constant (Kc) for given reactions.
  • Analyze the effects of changing system conditions on equilibrium positions.

Key Points

  1. Chemical Equilibrium Definition:
  2. Occurs when the forward and reverse reactions occur at the same rate in a closed system.
  3. The concentrations of reactants and products remain constant over time, signifying a dynamic balance.

  4. Equilibrium Constant (Kc):

  5. For a reaction (aA + bB \leftrightarrow cC + dD), the equilibrium constant expression is:
    [
    Kc = \frac{[C]^c [D]^d}{[A]^a [B]^b}
    ]
  6. The concentrations (([ ])) are in mol·dm(^{-3}).
  7. Pure solids and liquids are not included in the Kc expression 【4:3†source】.
  8. If ( Kc \gg 1 ), products are favored; if ( Kc \ll 1 ), reactants are favored.

  9. Factors Affecting Equilibrium:

  10. Concentration: Increasing the concentration of reactants/products shifts the equilibrium to oppose the change.
  11. Temperature: Increasing temperature favors the endothermic direction; decreasing temperature favors the exothermic direction.
  12. Pressure: Increasing pressure favors the direction with fewer gas molecules; decreasing pressure favors the direction with more gas molecules.
  13. Catalysts: Lower the activation energy, speeding up both forward and reverse reactions but do not change the Kc or the equilibrium position 【4:3†source】【4:7†source】【4:17†source】.

  14. Le Chatelier’s Principle:

  15. If a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change.

Real-World Applications

Example 1: The Haber Process

The Haber process synthesizes ammonia ( (NH_3) ) from nitrogen and hydrogen:
[ N_2(g) + 3H_2(g) \leftrightarrow 2NH_3(g) ]
Increasing Pressure: Shifts equilibrium towards ( NH_3 ) formation (fewer gas molecules).
Increasing Temperature: Shifts equilibrium towards ( N_2 ) and ( H_2 ) (endothermic reverse reaction)【4:4†source】【4:5†source】.

Example 2: Industrial Synthesis of Sulfuric Acid

Sulfur dioxide ( (SO_2) ) and oxygen ( (O_2) ) react to form sulfur trioxide ( (SO_3) ):
[ 2SO_2(g) + O_2(g) \leftrightarrow 2SO_3(g) ]
Increasing Pressure: Shifts equilibrium towards ( SO_3 ) (fewer gas molecules).
Catalysts: Increase reaction rate without affecting equilibrium position【4:8†source】.

Common Misconceptions and Errors

  1. Equilibrium does not imply equal concentrations of reactants and products.
  2. Catalysts do not change the position of equilibrium; they only speed up the rate at which equilibrium is reached.
  3. Equilibrium constants (Kc) are not affected by changes in concentration or pressure, only by temperature.

Strategies to avoid these errors:
– Carefully analyze the direction of the shift in equilibrium during changes.
– Remember that Kc is only altered by temperature changes.

Practice and Review

Practice Questions:

  1. Consider the reaction:
    [ H_2(g) + I_2(g) \leftrightarrow 2HI(g) ]
  2. If 0.4 mol of ( H_2 ) and 0.4 mol of ( I_2 ) react in a 1.5 dm( ^3 ) container at equilibrium, calculate ( [HI] ) (Kc = 45.9) 【4:3†source】【4:7†source】.

  3. For the reaction ( 2NO(g) + Br_2(g) \leftrightarrow 2NOBr(g) ):

  4. Calculate Kc if starting with 0.25 mol NO and 0.10 mol Br( _2 ) in 250 cm( ^3 ) container, with 0.20 mol NO at equilibrium 【4:4†source】.

Examination Tips:

  • Identify keywords: “shift” indicates a change in equilibrium position.
  • Use Le Chatelier’s Principle to predict shifts.
  • Always include units in equilibrium constant calculations.

Connections and Extensions

  • Interdisciplinary Links: Chemical equilibrium concepts are applicable in environmental science (atmospheric reactions), biology (respiratory gases), and industrial engineering (reaction optimizations).
  • Real-World Implications: Understanding equilibrium helps in pollution control by optimizing industrial reactions to minimize harmful emissions.

Summary and Quick Review

  • Chemical equilibrium is when forward and reverse reaction rates are equal.
  • The equilibrium constant (Kc) is only changed by temperature.
  • Le Chatelier’s Principle helps predict system shifts in response to changes.

Additional Resources

  • Online Articles:
  • Khan Academy (Chemical Equilibrium): Link
  • Videos:
  • YouTube Channel “Crash Course Chemistry” on Le Chatelier’s Principle: Link
  • Educational Platforms:
  • Coursera’s Chemistry Courses: Link

These resources provide further explanations and practice problems to deepen your understanding of chemical equilibrium.


Given the detailed content analysis and real-world scenarios, these revision notes are designed to provide comprehensive support to Grade 12 students preparing for exams in Chemical Equilibrium under the CAPS curriculum.

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