Kinetics

• Activation energy is the minimum energy required to start a reaction by the breaking of bonds.
• Reactions can only occur when collisions take place between particles having sufficient energy.
• Most collisions do not lead to a reaction because:

  1. particles don't have the activation energy
  2. particles may not be in the correct orientation
• Rate of reaction is defined as the change in concentration of a substance in a given time.
• The units of rate are mol dm-3 s-1
• On a concentration-time graph, the rate at a particular time is equal to the gradient at that time.
• Concentration:

  • Increasing the concentration increases the rate.
  • There are more particles in a given volume, so more frequent collisions, so more frequent successful collisions in a given time.
• rate = change in [] / time
• Pressure:

  • Increasing the pressure of gases increases the rate.
  • More particles in a given volume, so more frequent collisions, so more frequent successful collisions in a given time.
  • NOTE: increasing the pressure means either more moles in the same volume or the same moles in a smaller volume.
• Surface Area:

  • Increasing the surface area of solids (smaller particles) increases the rate.
  • More particles are exposed to the other reactant, so more frequent collisions, so more frequent successful collisions in a given time.
• Temperature:

  • Increasing the temperature increases the rate.
  • Increasing temperature gives all particles more energy so many more particles will have energy greater than or equal to Ea, so many more successful collisions per unit time.
• Catalysts:

  • A catalyst increases the rate of a reaction without being used up.
  • It provides an alternative reaction pathway with a lower activation energy, so more particles with energy greater than or equal to Ea, so more successful collisions per unit time.
• Continuous monitoring method:

  e.g. Small pieces of calcium carbonate react with HCl. CO2 is produced and is collected in a gas syringe. The volume of CO2 is recorded at regular intervals. A graph is plotted of CO2 produced against time.
• When the particles within a sample (ideal gas) collide, energy is exchanged between them so the particles have a spread of energies. The distribution of energies within the particles can be shown with the Maxwell-Boltzmann distribution.
• Key features of the Maxwell-Boltzmann distribution:

  • Starts at (0,0) as no molecules have zero energy.
  • The curve approaches the x-axis but never touches it (asymptotically) as there is no maximum energy a molecule can have.
  • Curve is skewed to the left because most probable energy/modal energy (Emp) is low.
  • Ea lies to the right (higher than the mean energy)
  • Area under curve represents total number of molecules in sample.
  • Shaded area represents number of molecules with energy greater than or equal to Ea.
• Features of Maxwell-Boltzmann with a temp increase:

  • area under curve stays the same as same number of molecules in sample.
  • curve lies to the right as all particles have more energy so Emp is higher.
  • shaded area is higher.
  • more collisions will have energy greater than or equal to Ea.
  • more collisions will be successful, so more successful collisions per unit time.
• In general, a small temp increase of 10 degrees causes rate to double. This large increase in rate is because increasing the temperature gives ALL particles more energy so many more particles have energy greater than or equal to Ea so many more collisions will be successful, so many more successful collisions per unit time.
• A catalyst is a substance which increases the rate of a chemical reaction without being changed in chemical composition or amount. It provides an alternative reaction pathway with a lower activation energy.