Temperature and Equilibrium

Introduction to Chemical Equilibrium

Definition and Importance

• Chemical Equilibrium: Chemical equilibrium is a condition where the concentrations of reactants and products remain unchanged over time, even as underlying microscopic reactions continue.
  • Significance: Essential for sustaining ecological equilibrium (e.g., balance of oxygen and carbon dioxide) and increasing industrial efficiency in processes such as ammonia production for fertilisers.

Dynamic Nature and Reaction Rates

• Microscopic Dynamics: Reactions at equilibrium persist microscopically. Imagine it akin to a balanced scale—maintaining equilibrium as particles consistently move.

  • Example: The dissolution and recrystallisation of salt in water illustrate this dynamic equilibrium process.

• Reaction Rate Balance: At equilibrium, the forward and reverse reaction rates are equal, resulting in no net change in concentration.

Equilibrium Constant ($K_{\text{eq}}$)

• Equilibrium Constant ($K_{\text{eq}}$): Indicates the equilibrium position, influenced solely by temperature.
• Mathematical Expression: $K_{\text{eq}} = \frac{[C]^c[D]^d}{[A]^a[B]^b}$
  • Each concentration is related to its corresponding ratio in the balanced chemical equation, as demonstrated in acid dissociation in vinegar.

Visualising Equilibrium

• Energy Profile Diagram: Illustrates changes in energy levels and the activation energy required to attain equilibrium.

  

• Concentration-Time Graphs: Demonstrate how reactant and product concentrations evolve over time to establish equilibrium.

  

Reversible Reaction Example

• Consider the reaction: 2SO₂ + O₂ ⇌ 2SO₃.

  

Key Points

• Equilibrium is dynamic, requiring a balance between forward and reverse reactions.
• Significance of $K_{\text{eq}}$: It determines the equilibrium position, critical for predicting reaction outcomes.

Le Chatelier's Principle in Temperature Effects

Introduction

• Le Chatelier's Principle: If a system at equilibrium experiences a change in conditions, it will adjust itself to partially counteract the change.
• Temperature is a crucial factor that influences the position of equilibrium.

The Role of Temperature in Chemical Systems

• Increasing temperature shifts equilibrium towards the direction that absorbs heat (endothermic).
• Decreasing temperature shifts it towards where heat is released (exothermic).

Distinguishing Reaction Types

• Endothermic Reaction:
  • Absorbs heat.
  • Example: The decomposition of calcium carbonate ($\text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2$).
• Exothermic Reaction:
  • Releases heat.
  • Example: The combustion of propane ($\text{C}_3\text{H}_8 + 5\text{O}_2 \rightarrow 3\text{CO}_2 + 4\text{H}_2\text{O}$).

Visual Representation of Temperature Effects

Energy Diagrams show changes in equilibrium:

• Exothermic Reaction Energy Diagram:

  • 
  • Takeaway: Increased temperature shifts equilibrium towards the reactants.

• Endothermic Reaction Energy Diagram:

  • 
  • Takeaway: Rising temperature shifts equilibrium towards the products.

Example Reaction

Reaction: $\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$

• Increase in Temperature:
  • Shifts equilibrium left (reverse).
• Decrease in Temperature:
  • Shifts equilibrium right (forward).

Temperature Effects on $K_{\text{eq}}$

Practical Context

• Haber Process:
  • Demonstrates the importance of $K_{\text{eq}}$ in optimising ammonia production.

Endothermic and Exothermic Reactions

• Endothermic Reactions:
  • An increase in temperature results in a rise in $K_{\text{eq}}$.

• Exothermic Reactions:
  • An increase in temperature leads to a decline in $K_{\text{eq}}$.

Introduction to Experimental Investigations

Purpose

• Comprehend: The influence of temperature on chemical equilibrium.

Cobalt Chloride Equilibrium Experiment

• Equation: $[\text{Co}(\text{H}_2\text{O})_6]^{2+} + 4\text{Cl}^- \rightleftharpoons \text{CoCl}_4^{2-} + 6\text{H}_2\text{O}$
• Temperature Impact:
  • Heating shifts equilibrium toward the blue CoCl₄²⁻.
  • Cooling shifts it toward the pink [Co(H₂O)₆]²⁺.

Nitrogen Dioxide/Dinitrogen Tetroxide Equilibrium

• Equation: $2\text{NO}_2 \rightleftharpoons \text{N}_2\text{O}_4$

Worked Example: When nitrogen dioxide gas (NO₂, brown colour) is cooled, it forms dinitrogen tetroxide (N₂O₄, colourless).

1. Initially at room temperature, the flask contains mostly NO₂ (brown colour).
2. When placed in an ice bath (0°C), the colour fades as N₂O₄ forms.
3. When removed and warmed to room temperature, the brown colour returns.

This demonstrates that the reaction 2NO₂ ⇌ N₂O₄ is exothermic (releases heat when forming N₂O₄), as cooling favours the forward reaction.

• Description:
  • Higher temperatures increase the concentration of brown NO₂.
  • Lower temperatures favour the formation of clear N₂O₄.

Iron(III) Thiocyanate Experiment

• Equation: $\text{Fe}^{3+} + \text{SCN}^- \rightleftharpoons \text{Fe(SCN)}^{2+}$
• Temperature Impact:
  • Increasing temperature favours the formation of red Fe(SCN)²⁺.

Safety Protocols and Equipment

• Protocols:
  • Always use protective gear and adhere to spill procedures.