 In this final kinetics video, we're going to get down to the nitty-gritty of exactly how things like concentration and temperature affect the rate of reaction. To do this, we're going to expand on collision theory, the theory that describes how collisions between molecules lead to reactions. In the last video, we talked about the facts that the rate of reaction actually depends on the chances of a successful collision occurring between reactant molecules. Now we're going to figure out why it is that the variables of surface area, concentration and pressure, temperature and catalysts affect those chances and thereby affect the reaction rate. Firstly, let's look at surface area, concentration and pressure. We're going to bundle these together because they essentially work in the same way. Take our hypothetical flask of reactants and their Maxwell-Boltzmann distribution. Have a look at the y-axis of this graph. This is the number of particles. Keeping all else the same, if we increase the surface area, concentration or pressure of one of the reactants, we're effectively increasing the number of reactant molecules per unit volume that are available to react. So we can redraw the distribution to show that there's more of everything like this. The shape of the distribution remains the same. Give or take my dodgy drawing. The proportion of particles with any particular speed doesn't change. It's just that there are more of them crowded together. Now, it's still the same reaction, so the activation energy remains in the same place. And if the activation energy is the same, then the proportion of particles over the activation energy also remains the same. But because we now have more particles in the same volume, the rate of all collisions increases, and so the rate of successful collisions must also increase. Just as an aside, we should make it explicit why increasing the surface area in a heterogeneous reaction has the same effect as increasing concentration in a homogeneous reaction. If the reactants are uniformly dispersed in a single homogeneous solution, the number of collisions per unit time depends on concentration and temperature. But if the reaction is heterogeneous, the reactants are in two different phases, and the reactant collisions can only occur at the interface between those phases. So the number of collisions per unit time is reduced relative to the homogeneous case, and so is the reaction rate. But if you increase the surface area, and you could do this by dividing one of your phases more finely, for example, like having calcium carbonate powder in acid as opposed to a single large chunk of calcium carbonate in acid. If you do this, then you increase the frequency of the collisions because there's more surface area, there's more reactant molecules exposed, and hence you increase the reaction rate. Car engines use surface area effects to increase reaction rates. By spraying the fuel into the engine cylinder in microscopic droplets rather than in, say, a single liquid stream, the fuel burns more rapidly. So to sum this up, increasing surface area concentration or pressure has the effect of increasing the number of collisions that each molecule undergoes. If the chance of a successful collision remains the same, but collisions happen more frequently, then successful collisions will also occur more frequently, and this means the reaction goes faster. To give you a really simple numerical example, imagine a situation in which molecules were undergoing 100 collisions per second with a 30% chance of being successful. That would mean that on average you would have 30 successful collisions per second. Now if you increase the concentration so that there are 200 collisions per second, you've still got a 30% chance of being successful, but that gives you, on average, 60 successful collisions per second. So we've effectively increased the rate of reaction.