The speed of a chemical reaction is based on the height of the activation energy. And we're going to say here that activation energy, which is abbreviated as \( e_a \) equals your transition state minus your reactant line. So remember, when we're looking at any particular energy diagram, the top of the curve is our transition state. And where the curve starts, that's where our reactants are initially. So we're going to say here that the higher the activation energy, then fewer reactive molecules have enough energy to convert into products. Remember that the difference in height between your transition state and your reactant line, that represents your activation energy.
So if we look at both of these, this would be our activation energy. So basically, we're going to say the higher your activation energy, then the slower the reaction. And we're going to say the smaller your activation energy then the faster the reaction. If we take a look here at these two reactions, remember, we just said that activation energy equals transition state minus your reactant line. So in this first curve, our transition state looks like it resides on this line here which is 90 minus your reactive line. Your reactive line here resides at 20. So the difference here is 70 kilojoules. So that's the minimum energy required for our reactants to traverse this curve and hopefully become products.
Over here, our transition state is 60 and our reactant line is still 20, so that's 40 kilojoules. Remember, we just said that the larger your activation energy, the slower your reaction. So we'd say that the first graph is our slower reaction. The second graph has a smaller activation energy, so this is our faster reaction.
Right? So just keep in mind that your energy of activation or activation barrier, this basically determines how fast your reaction will go. The higher your \( e_a \) is, the slower the reaction. The smaller your \( e_a \), the faster the reaction.