Catalysts

1. Catalysts How they affect the rate of reaction

2. A normal reaction:

3. A catalysed reaction

4. Analogy Suppose you have a mountain between two valleys so that the only way for people to get from one valley to the other is over the mountain. Only the most active people will manage to get from one valley to the other. Now suppose a tunnel is cut through the mountain. Many more people will now manage to get from one valley to the other by this easier route. You could say that the tunnel route has a lower activation energy than going over the mountain. But you haven't lowered the mountain! The tunnel has provided an alternative route but hasn't lowered the original one. The original mountain is still there, and some people will still choose to climb it. In the chemistry case, if particles collide with enough energy they can still react in exactly the same way as if the catalyst wasn't there. It is simply that the majority of particles will react via the easier catalysed route.

5. "A catalyst provides an alternative route for the reaction with a lower activation energy."

6. Industrial Examples of Catalysts Haber Process Uses finely divided iron to catalyse the production of ammonia from nitrogen and hydrogen. The two reacting gases absorb onto the active sites on the iron particles. The bonds within the reacting molecules are then weakened and new bonds form due to the close proximity. This proximity is what increases the rate of reaction compared to the two gasses just reacting in the air.

7. Uses of products made by the Haber process Fertiliser – further reactions are needed to make (NH2)2CO Explosives – ammonium nitrate Refrigeration and air conditioning Pharmaceuticals Paper and pulp making Cleaning reagents

8. Catalysts in the Environment Ozone Depletion Chlorine free radicals are formed from UV radiation reacting with chlorofluorocarbons – CFC’s Chlorine radicals react with ozone and break it down into oxygen gas. Cl· + O3 → ClO· + O2 ClO· + O· → Cl· + O2