 So far, we've been looking at experimental data for rates, analyzing how the rate changes as we change other factors. However, this doesn't tell us why the rate depends on factors like concentration, temperature, and so on. In the next few videos, we're going to go into this why by exploring the ideas of collision theory. To do this, we need to get right down to basics. If you shrank yourself to the size of a molecule and you were watching Adam's molecules undergo a reaction, what would you see? What is actually necessary at a molecular level for a reaction to occur? Well, for a reaction to occur, some or all of the bonds in the reactant molecules need to break so that the atoms can rearrange, create new bonds, and form the product molecules. So here are three conditions that are necessary for a reaction to occur. And they all have to do with the need to break bonds in the reactants. First, if two molecules are going to react, they need to be physically close to one another. And the way this happens with molecules is that they collide. Second, when they collide, the molecules have to collide in just the right way. What do we mean by just the right way? Well, here's an example. Nitrogen monoxide and ozone are able to react to produce nitrogen dioxide and oxygen. Essentially, the ozone is giving up one of its oxygen atoms to the nitrogen monoxide. So for a successful reaction, you're looking for the NO and the O3 to collide. But there are a number of different ways that this might occur. They're not all going to give us the result that we're looking for. In this first example, the oxygen of the NO bumps into the ozone. But these two oxygens can't make a bond, so the molecules bounce off each other and continue on unchanged. Or the nitrogen end of the NO could bump into the ozone. This looks promising. But it's bumped into the central oxygen in the ozone, and it's unlikely that this will break bonds to both of the other oxygens. So again, the molecules bounce off and continue on. However, if the nitrogen of the NO collides with an end oxygen on the ozone, then the oxygen can break off the ozone and join on to the NO, giving NO2. And then we have a successful reaction. So you can see that an NO and an O3 could collide many times before just by chance they happen upon the correct orientation. And it's that word chance that turns out to be the key to understanding reaction rates. Well, so far so good, but there is one more condition. Even if two molecules collide in exactly the right orientation, nothing will happen if there's not enough energy in the collision to cause the bonds in the reactant molecules to break. Two molecules approaching one another may have everything going for them, but if they're moving too slowly, all that will happen is they'll gently bounce off each other and continue on their way. So the last requirement is that the molecules be moving fast enough, or in other words, have sufficient kinetic energy for the collision to break some bonds and cause a reaction. And it's this link between the activation energy and the rate of reaction that we'll look at in the next video.