2.2 ENERGY PROFILES, DISTRIBUTION CURVES AND CATALYSTS

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Hi everyone!! Today we are going to finish off this topic. As you probably saw from the title we are aiming to draw energy profiles for exothermic and endothermic reactions. BUT there are a couple of things you need to know before we continue. And what I mean is few definitions.

The most important phrase that we have here is activation energy. Nothing will make sense if you don't know what activation energy is.

ACTIVATION ENERGY - the minimum energy required for the particles to collide and react with each other. Symbol is Ea.

As you kow from previous post, particles collide with each other and when they collide with enough energy, we call it successful collision which basically means they react. This 'enough energy' is the activation energy. Think of it as the energy needed to activate the paricles for a reaction. Of course activation energy will be different for every reaction.

You can show the activation energy on energy profiles (diagrams) but first of all you need to consider the two types of reactions that can happen. One is exothermic and the other is endothermic. (I drawn the two types of diagrams on the picture above).

Exothermic

In exothermic reaction energy is released. This means that the particles have more energy at the beginning than they have at the end. And this is why in the diagram the reactants have a line higher than the products.

The symbol for change in enthalpy shows the difference in energy. But we are more interested in the other symbol wich is our activation energy. The reactants need to 'gain' the activation energy then they react and products are created.

Endothermic

In endothermic reaction the energy is absorbed. This means that the paricles have less energy at the beginingn than they have at the end. And this is why in the diagram the line for reactants is lower at the begninng and higher at the end.

Looking at both diagrams I hope you can see that much more energy is needed to get the activation energy in endothermic reactions. This is why exothermic reactions are the most common type of reaction. Do you know all of the reactions that become warmer as they proceed. Yes, they are all exothermic reactions because we consider heat as energy.

Now, we will focus on the Boltzmann distribution curve.

This is all to do with gas molecules. Gases consist of molecules. They are moving constantly with different speeds (energies). As one hits the other, their speed changes. The one who got hit will gain kinetic energy and the one who hit will lose some kinetic energy.

The energy of these molecules is called molecular energy. So:

The Boltzmann distribution - distribution of molecular energies in a gas at a constant temperature.

Now, we know that all molecules have different energies or different speeds, but only a small part of these molecules reaches the activation energy (energy needed for them to collide successfuly and react). This is all shown on a Boltzmann distribution curve, right above in the picture.

HOWEVER, what happens if we increase the temperature?

Look at the graph and you can see that a bigger part of the molecules now reach the actiation energy. This means that as you increase the temperature, the molecules gain more kinetic energy, they collide with more energy and more of them have enough energy to react.

Some tips from me about the curves:

- never touch the x-axis

- areas under each of the curves (T1 and T2- with higher temperature) are the same. We still have the same amount of molecules but at higher temperature more of them reach the activation energy.

- the peak of T2 moves to the right.

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CATALYSTS

In this part we are going to define catalysts and I will show you how catalysts change both types of digrams that we lookaed at today.

Catalysts - increase the rate of reaction by providing an alternative route, without being used up.

What I mean by saying 'an alternative route'? This is the route with lower activation energy. So, particles need to gain less energy than without a catalyst.

1. The first diagram is with an exothermic reaction because as I said they are more common.

With orange colour I drew the activation energy for a particle without a catalyst and then I drew another line with blue colour showing the activation energy with a catalyst. It is clear to see that catalyst reduces the activation energy. This is how you would show the action of a catalyst on an exothermic reaction on a diagram.

2. The other diagram is showing the action of catalyst on a Boltzmann distribution curve. Again the activation energy is lower when we have a catalyst. This means more molecules have enough energy to react.

All of this overlaps with the previous post: catalyst increases the rate of reaction.

You just need to know a little bit more about catalysts.

There are two types of catalysts that you need to know for AS-level chemistry.

Homogenous - in the same phase as reactants.

e.g. concentrated sulfuric acid in making an ester (you will know more about this later on in this unit).

Heterogenous - in a different phase than reactants.

e.g. nickel in hydrogenation of hydrocarbons

e.g. iron in Haber process

I know that these examples mean nothing to you yet, but it will all become more clear later on in this unit. Just wait for it.

PS. Please remember, I am only a student, and as anyone, I can make mistakes. If you think you can see one, don't hesitate and comment (either here on on my youtube channel) Thank you!