Chemical kinetics and equilibrium

• Collision Theory


  • Reactant particles must collide for a chemical reaction to occur
  • Collide with the correct orientation
  • Collide with sufficient energy to overcome activation energy (EA)
• Activation energy (EA)


  Minimum amount of energy colliding particles must have for a specific chemical reaction to occur
• In general, faster reactions tend to have lower EAs
• EAs do not change with temperature
• Maxwell-Boltzmann distribution


  • Describes the distribution of speeds (kinetic energy) among the particles in a sample of gas at a given temperature
  • The total area beneath the curve is equal to the total number of particles in the sample

  
• As the temperature increases

  The particles gain kinetic energy and the curve flattens out
• Gas particles with lower molar masses such as hydrogen and helium have, on average, higher speeds than gas particles with higher molar masses such as oxygen and nitrogen
• Factors Affecting the Rate of Reaction

  • Temperature
  • Concentration/Pressure
  • Surface Area
  • Catalyst
• Rate of reaction


  • Change in concentration of a reactant or product per unit of time
  • Always a positive value
  • Measured in mol dm-3 s-1
• Measuring rates of reaction experimentally

  • Gas Production - Gas Syringe
  • Gas Production - Mass Loss
  • Time Taken to Form a Precipitate
  • Change in Concentration by Titration
  • Measuring the change in ion concentration
• The most appropriate way to measure the rate of reaction depends on the specific reaction
• Reversible reactions


  Reactions that can go in both the forward and backward direction, represented by the double-headed arrow symbol (⇌)
• Reversible reaction


  • N2 + 3H2 ⇌ 2NH3
• Forward reaction


  N2 + 3H2 → 2NH3
• Backward reaction


  2NH3 → N2 + 3H2
• Dynamic equilibrium


  Equilibrium where the rate of the forward reaction is equal to the rate of the reverse reaction, resulting in constant concentrations of products and reactants
• Dynamic equilibrium


  • Requires a closed system
  • No change in macroscopic properties
• Types of equilibria


  • Physical equilibria
  • Chemical equilibria
• Chemical equilibria


  • CO2(g) ⇌ CO2(aq)
  • CO2 + H2O ⇌ H2CO3 ⇌ HCO3- + H+
• Equilibrium constant, KC


  Describes the ratio of products to reactants at equilibrium
• Equilibrium constant has no units
• Changing the stoichiometry of KC

  Reverse the reaction: K2 = 1/K1

  Halve the coefficients: K2 = K1^(1/2)

  Double the coefficients: K2 = K1^2
• Reaction quotient, QC


  Measure of the relative amounts of reactants and products for a reaction not at equilibrium
• Le Chatelier's Principle


  When the conditions of a system at equilibrium change, the position of equilibrium shifts in the direction that tends to counteract the change
• Conditions that affect equilibrium

  • Concentration
  • Pressure
  • Temperature
  • Catalyst
• Changing conditions


  System response to shift the position of equilibrium
• Haber process


  Industrial process for producing ammonia from nitrogen and hydrogen
• Sources for Haber process


  • Nitrogen
  • Hydrogen
• Conditions for Haber process

  • Proportion of nitrogen to hydrogen
  • Temperature
  • Pressure
  • Catalyst
  • Ammonia removal from equilibrium mixture
• Factors affecting Haber process


  • Position of equilibrium
  • Rate of reaction
  • Economics
• Contact process


  Industrial process for producing sulfuric acid from sulfur dioxide and oxygen
• Sources for Contact process

  • Sulfur dioxide
  • Oxygen
• Conditions for Contact process

  • Temperature
  • Pressure
  • Catalyst
• Factors affecting Contact process

  • Position of equilibrium
  • Rate of reaction
  • Economics