3.2.2

• Rate of reaction


  The change in concentration or amount of a reactant or product per unit time
• Measuring rate of reaction


  1. Amount of reactant used or product made

  2. Divide by time
• Most collisions do not lead to reaction
• Conditions for reaction to occur

  • Particles must collide in the right direction
  • Particles must have sufficient kinetic energy
• Collision theory


  Theory of collisions required for reactions
• Activation energy


  Minimum energy required for a reaction to occur
• Energy profile diagram

  Shows changes in energy during a reaction
• Energy profile diagram


  • Reactants have stretching bonds and increasing kinetic energy
  • Activation energy is the minimum energy needed to break bonds
  • Products have lower energy than transition state
• Maxwell-Boltzmann distribution


  Shows the distribution of kinetic energies of gas particles
• Maxwell-Boltzmann distribution


  • Peak is the most likely energy
  • Mean is the average energy
  • Area under curve is total number of particles
  • Particles with energy above activation energy can react
• Increasing temperature


  Increases rate of reaction
• Effect of increasing temperature

  • Particles move faster and collide more often
  • More particles have energy above activation energy
• Increasing pressure


  Increases rate of reaction
• Reason for increased rate with pressure

  Particles are closer together, leading to more frequent collisions
• Increasing concentration


  Increases rate of reaction
• Reason for increased rate with concentration

  More particles present, leading to more frequent collisions
• Catalyst


  Substance that increases the rate of a reaction without being consumed
• How catalysts work


  • Provide alternative pathway with lower activation energy
  • Chemically unchanged at end of reaction
• Heterogeneous catalyst


  Catalyst in a different phase to the reactants
• Heterogeneous catalysts like iron are used in the Haber process to manufacture ammonia
• Porous catalysts


  • Increase the surface area of the catalyst
  • Allow reactions to happen much quicker
  • Increase the chance of collisions on the surface of the catalyst
• Reversible reactions


  Reactions that go forwards and backwards
• Reversible reaction progress


  1. Reactants used up quickly

  2. Concentration of reactants drops dramatically

  3. Products produced quickly

  4. Rate of forward reaction slows as reactants decrease

  5. Equilibrium reached when forward and backward rates are equal
• Homogeneous catalyst


  Catalyst is in the same phase as the reactants, normally in aqueous solutions
• Heterogeneous catalysts


  Reactants stick to the surface of the catalyst and then dissolve and move away, the catalyst is not used up
• Increasing concentration of a reactant or product


  Equilibrium shifts to reduce that concentration
• Increasing pressure


  Equilibrium shifts to the side with the fewest number of gas particles
• Homogeneous catalysts


  Form an intermediate by reacting with the reactants, then the product is reformed and the catalyst is used up
• Increasing temperature


  Equilibrium shifts in the endothermic direction
• Catalyst lowers the activation energy


  More particles now have enough energy to react
• Catalysts have no effect on the position of equilibrium, they only speed up the rate at which equilibrium is established
• Energy profile diagram


  Shows the activation energy is lower with a catalyst compared to without a catalyst
• Making ethanol


  • Reaction is exothermic going forward
  • Conditions: 60 atm pressure, 300°C, phosphoric acid catalyst
  • Lower temperature increases yield but decreases rate, so a compromise is reached
• Reasons catalysts are used in industry

  • Lower the temperature needed
  • Speed up the rate of reaction
  • Provide an alternative pathway
  • Change the properties of a product
• Environmental benefits of catalysts include lower temperatures and pressures, less energy required, and less waste produced
• Measuring reaction rates


  • Timing the formation of a precipitate
  • Measuring the mass of gas produced
  • Measuring the volume of gas produced
• High pressure


  Equilibrium shifts to the right
• Calculating reaction rate from a graph


  Use the gradient of the graph, which is the change in Y over the change in X
• For curved graphs, the reaction rate is calculated by drawing a tangent to the curve and finding the gradient of the tangent
• High pressurized equipment


  • Requires thicker, more robust vessels and pipes, which is expensive