Collision Theory
- Particles in a gas or liquid are constantly moving and colliding with each other. However, not all collisions result in a reaction.
- For a reaction to occur, two particles must collide with:
- Sufficient energy to overcome the activation energy
- The correct orientation
- Sufficient energy to overcome the activation energy
Maxwell-Boltzmann Distribution
Maxwell-Boltzmann Distribution
- The Maxwell-Boltzmann curve shows the distribution of molecular kinetic energies in a gas at a constant temperature
- In the Boltzmann distribution:
- The area under the curve is equal to the total number of molecules in the system
- The curve starts at the origin – no molecules have zero energy
- Only molecules with an energy greater than the activation energy, Ea, can react
- The curve’s peak represents the most probable energy that a molecule will have
- The average energy is to the right of, and slightly greater than, the most probable energy
Rate of Reaction
- The rate of reaction is the change in the concentration of a reactant or product in a given time.
- Factors affecting the rate of reaction:
- Temperature- increasing the temperature increases the kinetic energy of the molecules, leading to more frequent successful collisions and an increase in the rate of reaction
- Pressure – increasing the pressure of a gaseous reaction increases the number of gaseous molecules in a given volume, so molecules are closer together. This leads to more frequent successful collisions and an increase in the rate of reaction
- Concentration- increasing the concentration of an aqueous reactant increases the number of molecules in a given volume, so molecules are closer together. This leads to more frequent successful collisions and an increase in the rate of reaction
- Surface area- increasing the surface area of solid reactants increases the area over which a reaction can occur, leading to more frequent successful collisions and an increase in the rate of reaction
- Catalysts- adding a catalyst provides an alternative pathway that has a lower activation energy, leading to more frequent successful collisions and an increase in the rate of reaction
Catalysts
- A catalyst is a substance that increases the rate of a chemical reaction without being used up in the process.
- Catalysts work by reacting with the reactants to form an intermediate. Catalysts are then regenerated later in the reaction
- Homogeneous catalysts are in the same phase as the reactants
- Heterogenous catalysts are in a different phase to the reactants
- Heterogeneous catalysts work by:
- Adsorption- the reactants forming weak bonds with the atoms on the surface of the catalyst, holding the reactants in the correct position for them to react.
- Desorption- the products detach from the atoms on the surface of the catalyst
- As the products detach more reactants can be adsorbed, and the process is repeated
- The use of catalysts has economic, environmental and social benefits
- The demand for fossil fuels is lowered
- Emission of pollutants are lowered
- Reduced energy use means lower costs
The Effect of Temperature and Catalysts on The Maxwell-Boltzmann Distribution
- At higher temperatures, the kinetic energy of the molecules increases, so the molecules move faster
- A greater proportion of molecules have an energy greater than the activation energy, so more frequent successful collisions occur, increasing the rate of reaction.
- The area under the curve remains the same as the number of molecules in the system remains the same
- Catalysts lower the activation energy of a reaction by providing an alternative reaction route with a lower activation energy
- A greater proportion of molecules have an energy greater than the new, lower activation energy, and more frequent successful collisions occur, increasing the rate of reaction.
- The addition of the catalyst does not change the distribution of the molecular energies.
