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Chemical Equilibrium Reactions

  • Chemical reactions which can only go in one direction are known as irreversible reactions because the products formed cannot be combined to give the reactants under any conditions. Example of such reactions include:
    • 2KClO3(s) → 2KCl(s) + 3O2(g)
    • Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
  • Chemical reactions which can go in either direction depending on the conditions of the reaction are called reversible reactions. Examples include
    • CaCO3(s) + H2O(l) + CO2(g)  ↔ Ca(HCO3)2(aq)
    • H2O(s) ↔ H2O(l)
  • Chemical equilibrium reactions are reversible reactions carried out in closed containers which ensure that the products formed recombine to form the reactants at the same rate at which the reactants are form the products. Examples include
    • N2(g) + 3H2(g) ↔ 2NH3(g)
    • CaCO3(s) ↔ CaO(s) + CO2(g)
  • homogeneous equilibrium reaction is one in which the reactants and products are in the same physical state. The first example just above is an example of a homogeneous equilibrium reaction
  • heterogeneous equilibrium reaction is one in which the reactants and the products are not in the same physical states. The second example just above is an example of a heterogeneous reaction
Characteristics of a System in Chemical Equilibrium
  1. The rate of its forward reaction is exactly equal to the rate of its backward reaction
  2. The free energy change, ΔG, of an equilibrium system is equal to zero
  3. It has a dynamic nature because the forward and backward reactions never cease
  4. The concentrations of the reactants and products at equilibrium are constant at constant temperature, regardless of whether the equilibrium point is approached from the direction of pure reactants or pure products or a mixture of reactants or products
  5. Every chemical equilibrium system has a unique equilibrium constant, Keq, at constant temperature
Factors Affecting Equilibrium Reactions
  • The effect of changing conditions on any equilibrium can be predicted by Le Chatelier’s principle
  • The principle states that if a system in chemical equilibrium is disturbed by changing one of the factors maintaining equilibrium, the system will shift its equilibrium position in a direction that opposes or minimizes the imposed disturbance
  • In all cases, a disturbed equilibrium system always reduces the effect of an imposed disturbance by using one of its two reversible reactions to develop the exact opposite of the effect created by the imposed disturbance
  • The factors which can affect a system in chemical equilibrium include:
    1. Effect of Concentration
      • If the concentration of one substance is increased, the equilibrium position will shift in a direction that will rise up the substance whose concentration is increased
      • An increase in the concentration of a reactant always shifts the equilibrium position in favor of the products because product formation brings down the increased concentration of the reactants, the inverse is the case if there is an increase in the concentration of the products
      • If the concentration of one substance is removed from the system, the equilibrium position will move in the direction that favors the production of more of the substance being removed
    2. Effect of Pressure
      • This factor applies to only equilibrium systems in which at least one of the substances in the equilibrium system is in the gaseous state
      • It is specifically relevant to equilibrium systems in which the number of moles of gaseous reactants is not equal to the number of moles of gaseous products
      • In all cases, an increase in pressure (or decrease in volume of the reaction vessel) causes the equilibrium position to shift in favor of the side with fewer number of gaseous moles in the equilibrium system, the inverse is the case if there is a decrease in pressure (or an increase in the volume of the reaction vessel)
      • In general, if the pressure is increased, the equilibrium system will shift to reduce the pressure by reducing the number of gaseous particles present
    3. Effect of Temperature
      • The enthalpy change, ΔH, indicated for an equilibrium system is strictly for the forward reaction. ΔH for the backward reaction in an equilibrium system is always equal in magnitude but opposite in sign to that of the forward reaction
      • Thus, if the forward reaction is exothermic, then the backward reaction is necessarily endothermic, and vice versa
      • An increase in temperature will cause the equilibrium position to shift in the favor of the endothermic reaction (which absorbs heat) in the equilibrium system, the inverse would be the case if there was a decrease in temperature
    4. Effect of Catalyst
      • A catalyst has no effect on the equilibrium position or the yield of a chemical equilibrium reaction
      • A catalyst merely hastens the rate of the attainment of the equilibrium position by increasing the rates of the forward and backward reactions by equal amounts
Equilibrium Constant Keq
  • Every chemical equilibrium system has a unique equilibrium constant, Keq, which is constant at constant temperature
  • Keq is a quantitative measure of how far to the right the equilibrium reaction goes
  • If the value of Keq is very large, a large proportion of the reactants is converted to products
  • If the value of Keq is very small, no significant reaction takes place
  • Keq may be expressed in terms of concentration as Kcor in terms of partial pressures, Kp
  • For the reaction aA + bB ↔ cC + dD. Kc = [C]c [D]d / [A]a [B]b
  • If A, B, C, and D are all gases, Kp = PCc PDd / PAa PBb
  • If the equation of a chemical equilibrium reaction is reversed, the new equilibrium constant expression is the reciprocal of the former. Similarly, the new equilibrium constant is the inverse of the former
  • If a heterogeneous equilibrium reaction involves pure solids or pure fluids, the concentration of such pure solid or pure fluids are constant and so do not appear in the equilibrium constant expression. For example:
    • Fe(s) + 4H2O(g) ↔ Fe3O4(s) + 4H2(g). Kc = [H2]4 / [H2O]4 OR Kp = PH24 / PH2O4
  • The numerical value of an equilibrium constant, Kc or Kp must be found by experiment
  • The value of the equilibrium constant cannot be altered by pressure or concentration at a constant temperature
  • The only factor that can alter the value of an equilibrium constant is temperature
  • In general, an increase in temperature increases the equilibrium constant of an equilibrium reaction which is endothermic in its forward reaction.
  • If an equilibrium reaction is exothermic in its forward reaction, an increase in temperature decreases its equilibrium constant

Equilibrium Constant and Free Energy Change

  • The free energy change of a reaction is related to the equilibrium constant, K, by the equation
    • ΔGθ = -R T lnK
    • ΔGθ = -2.303R T log10K
    • ΔGθ is the standard free energy change of the system
    • K is the equilibrium constant
    • R is the universal gas constant
    • T is the temperature of the system in Kelvin
    • If K is very large (e.g. 106), ΔG is largely negative and the reaction is effectively complete
    • If the value of K lies between 1 and 106, ΔG is appreciably negative and the products predominate in the equilibrium system
    • Reactants predominate in the equilibrium system if K is lower than 1

Relationship between Kp and Kc

  • The equilibrium constant of chemical equilibrium reactions involving gases may be expressed as Kc or Kp
  • Kp and Kc for such reactions are related by the equation Kp = Kc(RT)Δn, where
    • Kp is the equilibrium constant in terms of partial pressure
    • Kc is the equilibrium constant in terms of concentration
    • R is the molar gas constant, 8.31 JK-1mol-1
    • T is the temperature of the reaction in Kelvin
    • Δn is the difference between the total number of moles of gaseous product and the total number of moles of gaseous reactants
  • For the reaction N2(g) + 3H2(g) ↔ 2NH3(g).  Kp = Kc(RT)-2
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