How does catalyst increase rate of reaction

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. It could be said that the tunnel route has a lower activation energy than going over the mountain, but the mountain itself is not lowered. The tunnel has provided an alternative route but has not lowered the original one.

The original mountain is still there, and some people still choose to climb it. In chemical terms, if particles collide with enough energy they can still react in exactly the same way as if the catalyst was not there; it is simply that the majority of particles will react via the easier catalyzed route. Jim Clark Chemguide. The importance of activation energy Collisions only result in a reaction if the particles collide with a certain minimum energy called the activation energy for the reaction.

The position of activation energy can be determined from a on a Maxwell-Boltzmann distribution: Only those particles represented by the area to the right of the activation energy will react when they collide. It can be represented on an energy level diagram. We will look at this in more detail in unit 2. The diagram shows that a catalyst provided a reaction pathway of lower activation energy. This makes more of the collisions successful at a given temperature.

You can mark the position of activation energy on a Maxwell-Boltzmann distribution to get a diagram like this:. Only those particles represented by the area to the right of the activation energy will react when they collide.

The great majority don't have enough energy, and will simply bounce apart. If there are very few particles with enough energy at any time, then the reaction will be slow. Important: I have already commented on this on the introductory page about collision theory and also on the page about the effect of temperature.

You mustn't get the idea that those particles in the blue area of the graph can never react. There are constant random collisions between the particles, and constant exchanges of energy between them.

Some particles will gain energy in random collisions, and others will lose energy. So a low energy particle could, an instant later, have gained enough energy from a collision that it could now react. And the opposite is true. A more energetic particle which didn't happen to collide successfully and produce a reaction, could find itself slowed down an instant later as a result of a collision.

Because of this constant exchange of energy, given time all the particles will react if the reacting proportions are right. To increase the rate of a reaction you need to increase the number of successful collisions. One possible way of doing this is to provide an alternative way for the reaction to happen which has a lower activation energy. As before, particles which don't have enough energy at a particular time will at some time in the future gain energy from random collisions, just as other particles will lose energy.

You mustn't get the idea that those particles in the blue area of the graph can never react - given time they will. Adding a catalyst has exactly this effect of shifting the activation energy. A catalyst provides an alternative route for the reaction.

That alternative route has a lower activation energy.