Chemical Equilibrium and Le Chatelier’s Principle

Posted on Oct 27, 2024 in Chemistry

Chemical Equilibrium

The chemical equilibrium is a reversible process where the rates of the forward and reverse reactions are equal, resulting in constant reactant and product concentrations over time. When all species are in the same phase, the equilibrium is homogeneous.

Homogeneous Equilibria: Law of Mass Action

In a homogeneous equilibrium, all chemical species are in the same phase. A general homogeneous equilibrium is represented as: aA + bB ⇌ cC + dD. At a given temperature, a constant ratio exists between the equilibrium concentrations of the substances. This ratio is the equilibrium constant (K_c).

A higher value of K_c or K_p indicates the reaction favors product formation.

Reaction Quotient

Q_c > K_c: The system is not at equilibrium. The product/reactant ratio is higher than at equilibrium. To reach equilibrium, Q_c must decrease by consuming products and forming reactants. The system shifts left.
Q_c = K_c: The system is at equilibrium.
Q_c < K_c: The system is not at equilibrium. The product/reactant ratio is lower than at equilibrium. To reach equilibrium, Q_c must increase by consuming reactants and forming products. The system shifts right.

Equilibrium Constant K_p

P_A = P_T
K_c = K_p / (RT)^-n
K_p = K_c (RT)ⁿ
If n = 0, K_c = K_p

Disturbing Equilibrium: Le Chatelier’s Principle

Changes in external factors (temperature, pressure, or concentration) affecting an equilibrium cause the system to adjust and establish a new equilibrium state, minimizing the effect of the change.

According to Le Chatelier’s principle, equilibrium composition can be altered by:

• Adding or removing a reactant or product:
  • Adding: Equilibrium shifts to consume the added excess.
  • Removing: Equilibrium shifts to form the removed substance.
• Changing pressure/volume (for gaseous reactions):
  • Compression: Increases total pressure and concentration, favors fewer gas moles.
  • Expansion: Decreases total pressure and concentration, favors more gas moles.
• Changing temperature (affects the value of K):
  • Increasing T: Favors the endothermic reaction.
  • Decreasing T: Favors the exothermic reaction.

Using a catalyst speeds up equilibrium attainment but doesn’t change K or the equilibrium composition.

Heterogeneous Equilibria

A heterogeneous equilibrium involves substances in different phases. Concentrations of pure solids, pure liquids, and the solvent are constant and omitted from the equilibrium expression.

• Gases are represented by partial pressures.
• Species in solution (aq) are represented by molar concentrations.
• Products appear in the numerator, reactants in the denominator, raised to their stoichiometric coefficients.

Free Energy and Equilibrium Constant

G = G⁰ + RT ln Q
G⁰ = – RT ln K
When G⁰ < 0, the reaction at equilibrium favors products.
When G⁰ > 0, the reaction at equilibrium favors reactants.

Van’t Hoff’s Equation

Van’t Hoff’s equation aligns with Le Chatelier’s principle.

• Exothermic reaction: K_p decreases with increasing temperature.
• Endothermic reaction: K_p increases with increasing temperature.

Degree of Dissociation

The degree of dissociation (α) is the fraction of a mole that dissociates in a reaction: α = (moles dissociated) / (initial moles)

Molecularity of a Reaction

Molecularity is the number of independent atoms or molecules involved in a reaction step (e.g., unimolecular, bimolecular).

Endothermic reactions absorb heat (positive enthalpy change).
Exothermic reactions release heat (negative enthalpy change).