THE COLLISION THEORY OF REACTION RATES This page describes the collision theory of reaction rates. It concentrates on the key things which decide whether a particular collision will result in a reaction - in particular, the energy of the collision, and whether or not the molecules hit each other the right way around (the orientation of the collision). The individual factors which affect the rate of a reaction (temperature, concentration, and so on) are discussed on separate pages. You can get at these via the rates of reaction menu - there is a link at the bottom of the page. We are going to look in detail at reactions which involve a collision between two species. Species: This is a useful term which covers any sort of particle you like - molecule, ion, or free radical Reactions where a single species falls apart in some way are slightly simpler because you won't be involved in worrying about the orientation of collisions. Reactions involving collisions between more than two species are going to be extremely uncommon (see below). Reactions involving collisions between two species It is pretty obvious that if you have a situation involving two species they can only react together if they come into contact with each other. They first have to collide, and then they may react. Why "may react"? It isn't enough for the two species to collide - they have to collide the right way around, and they have to collide with enough energy for bonds to break. (The chances of all this happening if your reaction needed a collision involving more than 2 particles are remote. All three (or more) particles would have to arrive at exactly the same point in space at the same time, with everything lined up exactly right, and having enough energy to react. That's not likely to happen very often!) The orientation of collision Consider a simple reaction involving a collision between two molecules - ethene, CH 2 =CH 2, and hydrogen chloride, HCl, for example. These react to give chloroethane. As a result of the collision between the two molecules, the double bond between the two carbons is converted into a single bond. A hydrogen atom gets attached to one of the carbons and a chlorine atom to the other. Diagram 1 The reaction can only happen if the hydrogen end of the H-Cl bond approaches the carbon-carbon double bond. Any other collision between the two molecules doesn't work. The two simply bounce
off each other. Of the collisions shown in the diagram, only collision 1 may possibly lead on to a reaction. If you haven't read the page about the mechanism of the reaction, you may wonder why collision 2 won't work as well. The double bond has a high concentration of negative charge around it due to the electrons in the bonds. The approaching chlorine atom is also slightly negative because it is more electronegative than hydrogen. The repulsion simply causes the molecules to bounce off each other. In any collision involving unsymmetrical species, you would expect that the way they hit each other will be important in deciding whether or not a reaction happens. Diagram 2 The energy of the collision Activation Energy Even if the species are orientated properly, you still won't get a reaction unless the particles collide with a certain minimum energy called the activation energy of the reaction. Activation energy is the minimum energy required before a reaction can occur. You can show this on an energy profile for the reaction. For a simple over-all exothermic reaction, the energy profile looks like this: If the particles collide with less energy than the activation energy, nothing important happens. They bounce apart. You can think of the activation energy as a barrier to the reaction. Only those collisions which have energies equal to or greater than the activation energy result in a reaction. Any chemical reaction results in the breaking of some bonds (needing energy) and the making of new ones (releasing energy). Obviously some bonds have to be broken before new ones can be made. Activation energy is involved in breaking some of the original bonds. Where collisions are relatively gentle, there isn't enough energy available to start the bond-breaking process, and so the particles don't react. The Maxwell-Boltzmann Distribution Diagram 3 Because of the key role of activation energy in deciding whether a collision will result in a reaction, it would obviously be useful to know what sort of proportion of the particles present have high enough energies to react when they collide. In any system, the particles present will have a very wide range of energies. For gases, this can be shown on a graph called the Maxwell-Boltzmann Distribution which is a plot of the number of particles having each particular energy.
The area under the curve is a measure of the total number of particles present. The Maxwell-Boltzmann Distribution and activation energy Diagram 4 Remember that for a reaction to happen, particles must collide with energies equal to or greater than the activation energy for the reaction. We can mark the activation energy on the Maxwell-Boltzmann distribution: Notice that the large majority of the particles don't have enough energy to react when they collide. To enable them to react we either have to change the shape of the curve, or move the activation energy further to the left. This is described on other pages. THE EFFECT OF CATALYSTS ON REACTION RATES This page describes and explains the way that adding a catalyst affects the rate of a reaction. It assumes that you are already familiar with basic ideas about the collision theory of reaction rates, and with the Maxwell-Boltzmann distribution of molecular energies in a gas. Note that this is only a preliminary look at catalysis as far as it affects rates of reaction. If you are looking for more detail, there is a separate section dealing with catalysts which you can access via a links at the bottom of the page. The facts What are catalysts? A catalyst is a substance which speeds up a reaction, but is chemically unchanged at the end of the reaction. When the reaction has finished, you would have exactly the same mass of catalyst as you had at the beginning. Some examples Some common examples which you may need for other parts of your syllabus include: reaction Decomposition of hydrogen peroxide catalyst manganese(iv) oxide, MnO 2
Nitration of benzene Manufacture of ammonia by the Haber Process Conversion of SO 2 into SO 3 during the Contact Process to make sulphuric acid Hydrogenation of a C=C double bond concentrated sulphuric acid iron vanadium(v) oxide, V 2 O 5 nickel The explanation The key importance of activation energy Collisions only result in a reaction if the particles collide with enough energy to get the reaction started. This minimum energy required is called the activation energy for the reaction. Diagram 5 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 Diagram 6 apart. Catalysts and activation energy 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. In other words, to move the activation energy on the graph like this: Adding a catalyst has exactly this effect on activation energy. A catalyst provides an alternative route for the reaction. That
alternative route has a lower activation energy. Showing this on an energy profile: A word of caution! Be very careful if you are asked about this in an exam. The correct form of words is "A catalyst provides an alternative route for the reaction with a lower activation energy." It does not "lower the activation energy of the reaction". There is a subtle difference between the two statements that is easily illustrated with a simple 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. Articles from http://www.chemguide.co.uk/physical/basicrates/catalyst.html http://www.chemguide.co.uk/physical/basicrates/introduction.html accessed on 2/23/09
AP Chemistry Chemical Kinetics Introduction Name Questions about The Collision Theory of Reaction Rates 1. First a bit of review. What are the four factors that can affect the rate of a chemical reaction? 2. In the middle of the first page, the statement is made that molecules first have to collide, and then they may react. What two things must be true about that collision in order for the molecules to react? 3. Reactions involving two molecules colliding in order to react are said to be bimolecular. Reactions involving three molecules colliding in order to react are called termolecular and are exceptionally rare. Why? 4. Sketch and label diagram 2. It s kind of an important diagram, so it s worth knowing it well. 5. Define the term activation energy. 6. What happens if particles collide without sufficient activation energy? 7. Which types of reactions require activation energy to proceed? 8. If we could calculate the average energy of a molecule in diagram 3, what would that value be called?
A Very Few Questions about The Effect of Catalysts on Reaction Rates 9. Copy and label diagram 5. 10. Adjust your just-drawn diagram to show what happens when a catalyst is added to a reaction. 11. The statement that a catalyst lowers the activation energy of a reaction is incorrect. What is the more correct but related statement about what a catalyst does? 12. Explain in your own words, of course the mountain/tunnel analogy used on the last page of the article and how it relates to catalysis. From neither article 13. Most reactions do not happen in one step. Instead, they happen in multiple steps. For example in the upper atmosphere, ozone is destroyed in this process. step 1 Cl + O 3 à ClO + O 2 step 2 ClO + O à Cl + O 2 What is the overall reaction taking place? 14. In the reaction from #13, what substance is the catalyst - the thing that is consumed then remade at the end? 15. What substance is an intermediate being made in one step then consumed in a later step?