The dynamic equilibrium between nitrogen dioxide (NO₂) and dinitrogen tetroxide (N₂O₄) provides an excellent visual demonstration of how chemical systems respond to environmental changes. This reversible reaction is particularly important in the study of Le Chatelier's principle.
The Equilibrium System
The equilibrium reaction between nitrogen dioxide and dinitrogen tetroxide can be represented by:
2NO2(g)⇌N2O4(g)ΔH=−58 kJ mol−1
This system is particularly useful for demonstration because:
NO₂ is brown in color
N₂O₄ is colorless
The reaction occurs readily at room temperature
Effects of Pressure and Volume Changes
When studying this equilibrium in a gas syringe, we can observe several key phenomena:
Increasing Pressure (Decreasing Volume)
When the syringe plunger is pushed in:
The volume decreases
Pressure increases
Concentration of both gases increases
The equilibrium shifts toward N₂O₄ formation
The brown color gradually fades
This shift occurs because:
The forward reaction produces fewer moles of gas (2 moles → 1 mole)
According to Le Chatelier's principle, the system reduces stress by favoring the reaction that produces fewer gas molecules
Decreasing Pressure (Increasing Volume)
When the syringe plunger is pulled out:
The volume increases
Pressure decreases
Concentration of both gases decreases
The equilibrium shifts toward NO₂ formation
The brown color intensifies
Temperature Effects
The reaction shows clear temperature dependence:
Heating the System
When heated in warm water:
The solution becomes browner
Equilibrium shifts toward NO₂ formation
The endothermic reverse reaction is favored
Cooling the System
When cooled in ice:
The solution becomes paler
Equilibrium shifts toward N₂O₄ formation
The exothermic forward reaction is favored
Summary
This equilibrium system demonstrates how:
Pressure changes affect systems with unequal moles of gaseous reactants and products
Temperature changes influence endothermic/exothermic reactions
Le Chatelier's principle predicts system behavior under changing conditions