Particles must collide with enough energy to break the bonds in order to react.
To increase the rate of reaction: increase temperature, increase concentration, increase pressure (gas only), increase surface area (solid only), use of a catalyst
Minimum energy required to start a reaction = activation energy, E
Activated complex = transition state
Provide a different reaction pathway with a lower activation energy.
Haber process = iron
Otswald process for making nitric acid = platinum/rhodium
Hardening of fats with hydrogen = nickel
Heterogeneous: where the catalyst is in different phase to the reactants, usually a solid catalyst and liquid or gaseous reactants.
Homogeneous: where the catalyst and reactants are in the same phase
Reduce levels of pollutants
Honeycomb of ceramic coated with platinum and rhodium
Shape = large surface area so a little expensive metal goes a long way
Gases pass over catalyst and react to form less harmful products
carbon monoxide + nitrogen oxides ---> nitrogen + carbon dioxide
hydrocarbons + nitrogen oxides ---> nitrogen + carbon dioxide + water
1. Gases form weak bonds with the metal atoms of the catalyst (adsorption) which holds the gases in the right position to react together. They react on the surface.
2. The products then break away from the metal atoms (desorption) which frees up room for more gases to react.