Chapter 10

• Reversible reaction

  After products form they can react to reform the original reactants
• Reversible reaction representation

  Two opposing half arrows are used: ⇌
• Dynamic equilibrium

  The reaction and products are constantly moving, the rate of forwards and backward reaction is the same, in a closed system with constant concentration of reactants and products

  
• In a closed system, the rate of H and I forming from hydrogen iodide is the same as the rate they are breaking into H and I

  
• Closed system

  None of the reactants or products escape from the reaction
• Open system

  Matter and energy can be lost to the surrounding
• Position of the equilibrium

  Refers to the relative amount of products and reactants in an equilibrium

  
• Left shift

  Concentration of reactants increases

  
• Right shift

  Concentration of products increases
• Le Chatelier's principle

  When a change is made to a system at equilibrium, the equilibrium shifts in a direction that counteracts or partially offsets that change

  
• Changes that shift the equilibrium

  • Adding reactant
  • Removing product
  • Increasing pressure
  • Increasing temperature

  
• Effect of catalyst

  They only cause reactions to reach equilibrium quicker, have no effect on the positioning of the equilibrium
• Investigating changes to the equilibrium position with concentration

  1. Adding H2SO4 increases rate of forward reaction, shifts equilibrium right

  2. Adding NaOH decreases rate of forward reaction, shifts equilibrium left

  
• Investigating changes to the equilibrium position with temperature

  1. Heating shifts equilibrium right (endothermic)

  2. Cooling shifts equilibrium left (exothermic)

  
• Equilibrium reactions

  Used in some stages of large-scale production of certain chemicals, important in the chemical industry
• Maximising ammonia yield in Haber process

  1. Increase pressure

  2. Decrease temperature

  3. Use catalyst

  
• Maximising sulphuric acid yield in Contact process

  1. Increase pressure

  2. Compromise temperature of 450°C

  3. Use vanadium(V) oxide catalyst

  
• Equilibrium expression and constant, Kc

  Kc is defined as the ratio of product concentrations raised to their stoichiometric coefficients to the reactant concentrations raised to their stoichiometric coefficients, solids are ignored

  
• Calculating Kc

  Use concentrations of reactants and products at equilibrium, Kc has units that depend on the equilibrium expression

  
• Interpreting Kc

  • If Kc is very large, equilibrium lies to the right (mostly products)
  • If Kc is very small, equilibrium lies to the left (mostly reactants)
  • If Kc is close to 1, similar concentrations of reactants and products
• Collision theory

  States that for a chemical reaction (collide successfully) to take place the particles need to collide with each other in the correct orientation and with enough energy

  
• Chemical reaction depend on

  1. How often the reactant particles collide

  2. How many successful collisions there are

  
• Ineffective collision

  When the particles collide in the wrong orientation and don't have enough energy to react so bounce off
• Increase in reaction rate

  1. More collisions happen when things bump into each other frequently

  2. When collisions occur more often, more particles get the energy needed to react

  3. This leads to a faster reaction rate

  
• Activation energy (Ea)

  The minimum amount of energy required for reactants particles to overcome for the reaction to take place (product formation)

  
• Exothermic reactions have activation energy lower than the reactants, endothermic reactions have activation energy higher than the reactants
• Factors that affect rate of reaction

  • Concentration (for liquids)/Pressure (for gases)
  • Surface area
  • Temperature
  • Catalyst
• Rate of reaction is the speed at which the chemical reaction is taking place, measured in mol dm-3 s-1 or mol dm-3 min-1

  
• Measuring rate from concentration-time graph

  1. Reactant concentration decreases

  2. Product concentration increases

  3. Rate of reaction is not constant and changes throughout the reaction

  
• Calculating rate of reaction at the start

  1. The concentration-time curve is almost linear at the start

  2. The curve becomes shallower with time as the rate decreases

  
• Catalyst

  Increases the rate of reaction by providing an alternative mechanism with lower activation energy, without being directly involved in the chemical reaction
• Types of catalysts

  • Homogeneous catalysts (same phase as reactants)
  • Heterogeneous catalysts (different phase to reactants)
• Heterogeneous catalysis

  1. Reactant adsorbs onto the surface of the catalyst

  2. Reaction takes place

  3. Product desorbs from the surface of the catalyst
• Benefits of catalysts

  • Speed up the rate of reaction, allowing lower temperature and pressure to be used
  • Enable different reaction routes, improving atom economy and reducing waste
  • Often enzymes, operating effectively at room temperature and pressure
• Homogeneous catalysts are in the same phase as the reactants and products, heterogeneous catalysts are in a different phase
• Catalysts decrease the activation energy of a chemical reaction, providing an alternative route with lower enthalpy change
• Maxwell-Boltzmann distribution

  A graph that shows the distribution of energies of particles at a certain temperature
• Only a small proportion of molecules in a sample have enough energy for an effective collision and chemical reaction to take place
• As temperature increases

  Particles gain more kinetic energy, causing more frequent collisions and a higher proportion of successful collisions with energy greater than the activation energy
• As temperature decreases

  The Boltzmann distribution curve becomes more peaked and shifted to the left, with fewer particles having enough energy for successful collisions