Kinetics is the study of the factors which affect the rate of a chemical reaction
- For a reaction to take place between two particles they must collide with sufficient energy to break bonds
- These reactions happen at certain parts of the molecule which means the orientation of the molecule also has to be correct.
Factors affecting rate of reaction
- Change in temperature - particles have more kinetic energy
- Increase in concentration of solution - more particles in a given volume at any given time leads to more collisions
- Increase in pressure (gas) - more particles in a given volume
- Increase in surface area of solid reactants - Greater total surface area means more particles are available to collide. More sites for reaction
- Use a catalyst - provides an alternative pathway for which there is a lower activation energy. Doesn't get used up during the course of the reaction
Collision Theory 2
Activation Energy, Ea - The minimum energy that a particle needs in order to react.
- At the transition state bonds are being both formed and broken.
- In an exothermic reaction energy of products is < reactants
Maxwell-Boltzmann Distribution shows the distribution of energy amongst particles
- Peak indicates most probable energy (Emp)
- Slight to the right of this is the mean energy
The graph shows that:
- No particles have zero energy
- Most particles have intermediate energies (around peak)
The line Ea inducates the activation energy and the area below the curve beyond this line indicates the no. of particles with E>Ea
- If the temperature is increased the peak (Emp) moves to the right and down but area underneath curve should still be the same. Area beyond Ea line will be greater.
- Catalysts provide an alternative pathway for which there is a lower Ea (smaller amount of energy needed to start reaction). With cataylst Ea line moves back.
With a catalyst the transition state for a chemical reaction is lower.
- Catalysts do not affect the enthalpy change of reactions or position of equilibrium
Heterogenous catalyst - The catalyst is in a different phase (state) to the reactants. Usually solid catalyst and liquid/ gas reactants
Homogeneous catalyst - The catalyst is the same phase as the reactants
Catalysts for well-known reaction
- Haber process - Iron
- Ostwald process for making HNO3 - Platinum
- Cracking hydrocarbons - Aluminium oxide/ Silicon oxide/ zeolite
- Honeycomb shape - provides a large surface area for reaction to take place
- Coated in Rhodium/Platinum (catalysts)
- Reactions: CO + NOx ---> N2 + CO2
- Metal forms weak bonds with reactants to hold in place - ADSORPTION
- Products then break away from metal - DESORPTION