Hi. It's Mr. Andersen and this is chemistry essentials video 39. And it's activation energy.
Imagine if I try to roll a ball up a hill. If I get enough energy into that system the
ball is going to roll down the other side. In other words there was a certain amount
of energy that that ball had at the beginning. We'll call that potential energy. And then
we've lost some of that potential energy. But the ball just didn't go there on its own.
I had to put a certain amount of energy into it. And in chemistry we call that activation
energy. Now it's not only balls rolling up a hill and down the other side. It's chemical
reactions. You can think of this as the reactants on this side. And this as products on the
other side. And so we're only netting this amount of energy between here and here. But
there's energy that you have to put into the system. Let's say we have an endergonic reaction
like this. One that requires energy. Well I'm going to put more energy in. I'm going
to get some of that back. If we were to look at the energy change then everything above
that is going to be the activation energy. What if I don't have enough activation energy?
Well I can put some energy in. But if I don't have enough, then it's going to go back to
where it was before. And so activation energy is how much energy you have to put into a
system to have a chemical reaction occur. And it's based on collisions between molecules.
Most of those are unsuccessful collisions. And that's because they maybe don't have enough
energy. Or they don't have proper orientation. But if they are successful, which again is
incredibly rare, then we can have chemical reactions fire off. In a unimolecular reaction
that's mostly going to be interactions between the molecule and the solvent or background
molecules. But if it's a bimolecular or termolecular reaction it's going to be with other molecules.
Or reactions with other molecules. In either case, it's going to be based on the Maxwell-Boltzmann
distribution, how often this is going to occur. And again that's where we can determine activation
energy. And so let's say we have two molecules. We'll call this A and B. And I were to launch
A at B, what we'll find is that if I don't give it enough energy then there's not going
to be a chemical reaction in this one simple collision. In other words you require a certain
amount of energy for that reaction to occur. So it's kind of based on temperature. If I
increase temperature I can have more reactions. Also if I fire it off center like that, then
they're not lined up properly and there's not going to be a reaction as well. And so
with each of these collisions between all the molecules you have to make sure that you
have proper energy and you also have proper orientation. But if you do, then we can fire
off that reaction. In other words we've exceeded that activation energy. And so this could
be in a unimolecular, when we just have one molecule turning into products. It could be
bimolecular where they're running into each other. Or it could even be termolecular. We
have three things coming together at once. And so all of these are based on the number
of collisions that we could possible have. And so the Maxwell-Boltzmann distribution
explains how much energy we have to put into a system for that chemical reaction to actually
occur. And so the curve looks like this. And so let me explain it to you. If you think
of this curve representing all the molecules that have a potential to interact, well we'd
say the mode is going to be right here. This is the number that have this amount of energy.
And this is the number of particles that we have. If we look on the left side right down
here, there are going to be no particles who have no energy. And then as we move to the
right side the mean is going to be a little bit to the right of that. Because there's
going to be a lot of particles way out here that have a huge amount of energy. Now there's
not many of those particles out there, but there's a lot of particles out there that
have a quite a bit of energy. And so if we look at that area under the curve, that represents
all the molecules that we have. Or have the potential to have interactions with. And so
we're going to set E sub A right here. That's going to be that activation energy. And so
all of these out here have the proper amount of energy to actually have a chemical reaction
occur. And so what happens if we increase the temperature? Well if we increase the temperature,
we just get a different curve. It's going to be the same number of particles, but it's
shifting everything to the right. Notice what happened to the activation energy. The activation
energy still is the same. It still requires a certain amount of energy for that reaction
to occur. But by shifting it higher temperature, we have more particles that have more energy.
And so that area out here is going to be greater. And so again, did you learn to explain the
difference between success and unsuccessful reactions? Again, it's based on the energy
they have and the orientation they have. And we can use the Maxwell-Boltzmann distribution
to figure that out. And I hope that was helpful.