Reaction Rates and Catalysts

1. Particles in liquids and gases are always moving and colliding with each other.They don't react every time though - only when the conditions are right. A reaction won't take place between two particles unless-

• They collide in the right direction. They need to be facing each other the right way.
• They collide with at least a certain minimum amount of kinetic (movement) energy.

This stuff's called Collision Theory.

2. The minimum amount of knietic energy particles need to react is called the activation energy. The particles need this much to break the bonds to start the reaction.

3. Reactions with low activation energies often happen pretty easily. But reactions with high activation energies don't. You need to give the particles extra energy by heating them.

To make this a bit clearer, here's a enthalpy profile diagram on the next card:

1. Reactants: Here the bonds within each particle are being stretched.

2. Peak of activation energy curve: If the particles have enough energy, the bonds will break.

3. Ea: This is the energy barrier that the particles have to overcome.

4.  Products: The seperate bits from each particle can't exist by themselves - so they form new bonds and release energy.

Imagne looking down on Oxford Street when it's teeming with people. You'll see some people ambling along slowly, some hurrying quicly, but most of them will be walking with a moderate speed. It's the same with the molecules in a gas. Some don't have much kinetic energy and move slowly. Others have loads of kinetic energy and whizz along. But most molecules are somewhere in between.

If you plot a graph of the numbers of molecules in a gas with different kinetic energies you get a Maxwell- Boltzmann distribution. It looks like this-

If you increase the temperature, the particles will on average have more kinetic energy and will move faster. So a greater proportion of molecules will have at least the activation energy and be able to react. This changes the shape of the Maxwell-Boltzmann distribution curve- it pushes to the right:

Because the molecules are flying about faster, they'll collide more often. This is another reason why increasing the temperature makes a reaction faster.

If you increase the concentration of reactants in a solution, the particles will on average be closer together.

If they're closer together, they'll collide more often. If there are more collisions, they'll have more chances to react.

If the reaction involves gases, increasing the pressure of the gases works in just the same way.

You can use catalysts to make chemical reactions happen faster. Learn this definition:

A catalyst increases the rate of a reaction by providing an alternative reaction pathway with a lower activation energy. The catalyst is hemically unchanged at the end of the reaction.

Catalysts are great. They don't get used up in the reactions, so you only need a tiny bit of catalyst to catalyse a huge amount of stuff. They do take part in reactions, but they're remade at the end.

Catalysts are very fussy about which reactions they catalyse. Many will usually work on a single reaction.

Catalysts save heaps of money in industrial processes.

If you look at an enthalpy profile together with a Maxwell-Boltzmann Distribution, you can see why catalysts work:

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