Equilibrium

• Equilibrium is a state in which the reaction mixture contains reactants and products whose concentration remains constant as long as temperature and pressure are constant
• Equilibrium involves two opposing processes and can be physical or chemical
• Physical Equilibrium applies to physical changes, while Chemical Equilibrium applies to chemical changes
• At equilibrium, the rate of two opposing processes are equal
• Properties like temperature, pressure, and concentration of the system do not change with time at equilibrium
• Common equilibrium involving physical processes include phase transformations like Solid-Liquid, Liquid-Gas, and Solid-Gas equilibriums
• In Solid-Liquid Equilibrium, the rate of melting of ice is equal to the rate of freezing of water at the normal freezing point
• In Liquid-Vapour Equilibrium, the rate of evaporation is equal to the rate of condensation, giving the equilibrium vapour pressure of water
• Solid-Vapour Equilibrium is observed in substances like ammonium chloride and iodine that undergo sublimation upon heating
• Equilibrium involving dissolution of solids or gases in liquids includes Solids in liquid equilibrium and Gas solution equilibrium
• In Solids in liquid equilibrium, the rate of dissolution is equal to the rate of precipitation, leading to a saturated solution
• Gas solution equilibrium follows Henry's law, where the mass of a gas dissolved in a solvent is directly proportional to the pressure of the gas above the solvent
• General characteristics of equilibrium involving physical processes:

  • An observable property of the system becomes constant at equilibrium
  • Equilibrium is possible in a closed container at a given temperature
  • Equilibrium is dynamic
  • Concentrations of substances become constant at equilibrium
  • An expression involving concentrations of substances is constant at constant temperature
• Equilibrium in chemical systems involves reversible chemical reactions carried out in closed containers
• Equilibrium constant represents the extent to which a process proceeds before equilibrium is attained
• Reversible reactions can take place in both the forward and backward directions under the same conditions
• Irreversible reactions cannot proceed in the backward direction even in closed containers
• Irreversible reactions are defined as reactions where the products do not react to give back the reactants under the same conditions
• Examples of reversible reactions:

  • 3 KClO3 (s) → 2 KCl (s) + 3 O2 (g)
  • AgNO3 (aq) + NaCl (aq) → NaNO3 (aq) + AgCl (s)
  • 2 Mg (s) + O2 (g) → 2 MgO (s)
• Equilibrium in heterogeneous systems:

  • Reactants and products are in different phases
  • Example: 3 CaCO3 (s) → CaO (s) + CO2 (g)
  • Equilibrium is reached when the rate of decomposition of CaCO3 is equal to the rate of formation of CaCO3
• Equilibrium in homogeneous systems:

  • Reactants and products are in the same phase
  • Example: 2 HI (g) ↔ H2 (g) + I2 (g)
  • Equilibrium is reached when the rate of formation of HI is equal to the rate of decomposition of HI
• Dynamic equilibrium:

  • Molar concentrations of reactants and products remain constant at equilibrium
  • Rate of forward reaction equals the rate of backward reaction
• Law of mass action:

  • Rate of a chemical reaction is directly proportional to the product of molar concentrations of reactants raised to their stoichiometric coefficients
• Law of chemical equilibrium:

  • At equilibrium, the ratio of the product of molar concentrations of products to reactants raised to their stoichiometry coefficients is the equilibrium constant
• Relation between equilibrium constants for molar concentrations and partial pressures in gaseous systems:

  • For ideal gases, the equilibrium constant using partial pressures is related to the equilibrium constant using molar concentrations by the ideal gas law equation
• Equilibrium constant when partial pressure is considered: pK = cK
• Equilibrium constant when molar concentration is considered: cK =
• Universal gas constant: R
• Temperature in Kelvin scale: T
• Difference between the gaseous products and reactants: nΔ
• Homogenous equilibrium examples:

  • H2 + I2 ⇌ 2HI
  • NH3 + H2O ⇌ NH4+ + OH-
• Heterogeneous equilibrium:

  • Molar concentration of pure solids and pure liquid is taken as 1
• Units of equilibrium constant:

  • For general reaction aA + bB ⇌ cC + dD:
  • If nΔ = 0, cK or pK does not have any unit
• Application of equilibrium constant:

  • Predict the extent of a reaction based on its magnitude
  • Predict the direction of a reaction
  • Calculate equilibrium concentrations
• Units of equilibrium constant:

  • For the reaction H2 + I2 ⇌ 2HI:
  • 1 mol atm or bar c pK =
• Factors affecting equilibria:

  • Change of concentration of any reactant or product
  • Change of temperature of the system
  • Change of pressure of the system
  • Addition of catalyst
  • Addition of inert gas
• When 2 H or 2 I is added to the reaction mixture at equilibrium, the equilibrium will shift in the direction where it is used up, leading to the formation of the forward reaction until a new equilibrium is reached
• When HI is added to the reaction mixture at equilibrium, the reaction shifts backward until a new equilibrium is established
• Pressure has a significant effect on reactions involving gases:

  • If the number of product moles is less than reactant moles, increasing pressure shifts the equilibrium in the forward direction
  • If the number of product moles is more than reactant moles, increasing pressure leads to the reverse reaction
  • If the number of reactant and product moles are the same, changes in pressure do not affect the equilibrium
• Temperature changes affect the equilibrium constant:

  • For exothermic reactions with a negative enthalpy change, increasing temperature shifts the equilibrium to the left and decreasing temperature shifts it to the right