 Let's take a hypothetical reaction where A and B react to give C and D and let's assume all of them are gases. So for this reaction what we want to know is what happens to the equilibrium when a catalyst is introduced? Does the equilibrium shift towards the reactants or the products or does it not change? So first we need to consider the role of the catalyst in the reaction. Now when we introduce a catalyst it is important to note that a catalyst is not permanently consumed in the reaction. So if we add a small amount of catalyst at the start of the reaction we'll get the same amount at the end of the reaction as well. And the other thing is when we introduce a catalyst there is no change in the other factors like volume, temperature, pressure or number of moles. So keeping these two points in mind if we write down the equation for equilibrium constant for this reaction it will look something like this which is the partial pressure of C times the partial pressure of D which are the products divided by the partial pressure of B times the partial pressure of A which are the reactants. And because we know that the catalyst is not consumed so the concentration of the catalyst or the partial pressure of the catalyst is not going to factor in here and the addition of a catalyst is not changing any of these factors and so we know that the partial pressures will not change. So if the partial pressures are not going to change that means the equilibrium constant is also not going to change. So we can say that a catalyst does not affect the equilibrium or in other words the catalyst is not going to change how much of the reactants or how much of the products will be present at the equilibrium. But then what is the catalyst doing? The catalyst is actually speeding up this reaction. So we can see that it affects the rate of the reaction that is how fast we get from the initial condition to the equilibrium but it is not changing the equilibrium constant or the position of equilibrium itself and a simple way to understand this difference is to think of this example. So let's say you and your friend are on the ground floor of this building and both of you want to get to the top floor which is marked by point B and one of you decides to take the staircase while the other person takes the lift and hopefully this is not a slow lift and so you know that taking the lift is faster than taking the stairs. So whoever will take the lift will get to point B quickly but the point is this top floor is not going to change depending on whether you take the lift or you take the stairs. Both of you are going to get at the same point but it's just that taking the lift gets you faster from A to B. That is what adding a catalyst does. It takes you faster to the destination but the destination itself does not change. Now let's look at the role of a catalyst using an energy diagram. We have this energy diagram here which is a plot of energy on the y-axis and the progress of the reaction on the x-axis. So as we go from the reactant side to the product side there is this barrier that has to be overcome or you can think that to get from the reactants to the products this much energy has to be provided but what happens when we introduce a catalyst. So what a catalyst does is that it stabilizes the reaction intermediate. So if you think of the energy at this point as that of the reaction intermediate what the catalyst does is it finds an alternative path for the reaction to proceed and with this path the activation energy is lower. So let's say before using the catalyst the activation energy for the forward reaction was this ER which is this distance and on using the catalyst the activation energy has now reduced to this EC and a very important point to note here is that the catalyst is not lowering the activation energy. The activation energy for the original reaction is still the same that is this ER. What the catalyst is doing is it is providing an alternative path and the energy for this path is lower. So the catalyst is not changing the original activation energy and that is an important distinction and the other thing is you can see here how a catalyst speeds up not just the forward reaction but also the reverse reaction for the forward reaction we can see that this EC is less than this ER. So as we move from reactants to products this barrier is reduced but for the reverse reaction when we go from the product side to the reactants even in that case the activation energy is lowered. So you can see here how for the reverse reaction without the catalyst the activation barrier is the sum of both of these ER plus E0 which is the total barrier for this reverse reaction which I am going to represent here by this EB and when you introduce a catalyst the activation barrier now becomes EC plus E0 which is a smaller value that I have denoted by this EB dash. So the point here is that the catalyst is not only speeding up the forward reaction it is also speeding up the reverse reaction by the same amount.