2.2 Le Chatelier's Principle and factors affecting equilibrium

Chemical equilibrium is a dynamic state where forward and reverse reactions occur at equal rates. Le Chatelier's Principle explains how systems respond to disturbances, shifting to counteract changes in concentration, pressure, or temperature.

Gibbs free energy connects thermodynamics to equilibrium. It determines reaction spontaneity and direction, linking to the equilibrium constant. Understanding these concepts helps predict and control chemical reactions in various applications.

Le Chatelier's Principle and Equilibrium

Le Chatelier's principle and equilibrium shifts

• Le Chatelier's Principle states that when a system at equilibrium is disturbed by a change in concentration, pressure, volume, or temperature, the system will shift its equilibrium position to counteract the disturbance and re-establish equilibrium (Haber process for ammonia synthesis)
• Changes in concentration affect equilibrium by shifting it towards the side that will reduce the concentration change
  • Increasing reactant concentration shifts equilibrium towards products to consume the added reactants (adding more CO in the water-gas shift reaction)
  • Decreasing reactant concentration shifts equilibrium towards reactants to replenish the consumed reactants (removing H2 in the Haber process)
  • Increasing product concentration shifts equilibrium towards reactants to consume the added products (adding more CO2 in the water-gas shift reaction)
  • Decreasing product concentration shifts equilibrium towards products to replenish the consumed products (removing NH3 in the Haber process)
• Changes in pressure or volume affect equilibrium in gaseous systems by shifting it towards the side with fewer moles of gas to minimize pressure change
  • Increasing pressure (decreasing volume) shifts equilibrium towards the side with fewer gas molecules (N2 + 3H2 ⇌ 2NH3)
  • Decreasing pressure (increasing volume) shifts equilibrium towards the side with more gas molecules (2SO2 + O2 ⇌ 2SO3)
• Changes in temperature affect equilibrium depending on whether the reaction is exothermic or endothermic
  • Increasing temperature shifts equilibrium towards the endothermic direction to absorb the added heat (N2 + O2 ⇌ 2NO)
  • Decreasing temperature shifts equilibrium towards the exothermic direction to release heat (2NO2 ⇌ N2O4)

Catalysts in reaction rates

• Catalysts accelerate the rate of a reaction by lowering the activation energy barrier without being consumed in the reaction (enzymes in biological systems, catalytic converters in automobiles)
• Catalysts increase the rate of both forward and reverse reactions equally, thus not altering the equilibrium constant ($K$) or the equilibrium position (Haber process using iron catalyst)
• Catalysts help the system reach equilibrium faster without shifting the equilibrium towards reactants or products (catalytic hydrogenation of vegetable oils)

Reactant and product effects on equilibrium

• Adding reactants increases their concentration, shifting the equilibrium towards products to counteract the disturbance (adding more CO in the water-gas shift reaction)
• Removing reactants decreases their concentration, shifting the equilibrium towards reactants to replenish the consumed species (removing H2 in the Haber process)
• Adding products increases their concentration, shifting the equilibrium towards reactants to counteract the disturbance (adding more CO2 in the water-gas shift reaction)
• Removing products decreases their concentration, shifting the equilibrium towards products to replenish the consumed species (removing NH3 in the Haber process)

Gibbs Free Energy and Equilibrium

Gibbs free energy and equilibrium position

• Gibbs free energy ($\Delta G$) is a thermodynamic quantity that determines the spontaneity and direction of a reaction at constant temperature and pressure
• The relationship between $\Delta G$ and the equilibrium constant ($K$) is given by the equation: $\Delta G = -RT \ln K$, where $R$ is the gas constant and $T$ is the absolute temperature
• $\Delta G < 0$ indicates a spontaneous reaction, with equilibrium favoring product formation (formation of rust, Fe2O3)
• $\Delta G > 0$ indicates a non-spontaneous reaction, with equilibrium favoring reactant formation (decomposition of water into H2 and O2)
• $\Delta G = 0$ indicates a system at equilibrium, with constant concentrations of reactants and products (saturated salt solution)
• The magnitude of $\Delta G$ determines the extent of the reaction: a more negative $\Delta G$ results in a larger equilibrium constant and a greater proportion of products at equilibrium (formation of NaCl from Na and Cl2)