Chemical Reaction Catalyst: Rates & Overview

Instructor: Nicola McDougal

Nicky has taught a variety of chemistry courses at college level. Nicky has a PhD in Physical Chemistry.

Catalysts are common in both industrial processes and in the human body. This lesson will explain how they work and why we need them. A quiz will test our knowledge.

Definition

In chemistry, a catalyst is any substance that speeds up a chemical reaction without itself being consumed in the reaction. Enzymes are naturally occurring catalysts responsible for many essential biochemical reactions.

What is a Reaction Rate?

Before we can think about catalysts we need to remind ourselves about the rates of reaction or reaction rate. Every chemical reaction has a reaction rate. This is the time it takes to get from reactants to products. At a very fundamental level, we think about chemical reactions as being the result of particle collisions.

Collisions only result in a chemical reaction if the reactant particles collide with a certain minimum energy. This energy is called the activation energy for the reaction. If a collision possesses an energy greater than the activation energy, that collision can result in a reaction. If a collision has an energy less than the activation energy, the molecules will bounce apart unchanged. The amount of activation energy needed will vary depending on the reaction. Normally, the lower the activation energy, the faster the reaction.

The diagram below shows a chemical reaction going from reactants on the left to products on the right. Energy is shown on the vertical axis. You will see that the energy goes up in the middle, and this is the activation energy required for this reaction. The reacting hydrogen and chloride molecules have to collide with enough energy to get over this hump and form hydrogen chloride molecules.

The reaction pathway. The Activation Energy is the energy needed to get from reactants to products
Diagram of the activation energy of a reaction

How Catalysts Work

To increase the rate of a reaction, you need to increase the number of successful collisions (collisions with sufficient energy). One possible way to do this is to provide an alternative way for the reaction to happen, which has a lower activation energy. You are still starting with the same reactants and finishing with the same products, but the reaction is going a different route to get there.

I often think of it like going on a journey. Say I want to go from my home town in Baltimore to Washington DC. The most important question is how long will it take to get there? The journey time is my reaction pathway. On one hand, I can go on the yellow tour bus travelling through all the towns on the way down to DC. This takes almost 2 hours. On the other hand, I could take the blue express bus. This time the bus goes a more direct route and I get there in just 1 hour and 5 minutes. I am still going from Baltimore and ending up in DC but the journey time is significantly less. The route the bus takes affects how fast I get there. Energy is also saved on the express bus because it goes a more direct route.

This is exactly the same in a chemical reaction. Adding a catalyst provides an alternative route for the reaction. That alternative route has a lower activation energy. This is shown on the diagram as a red line.

The reaction pathway this time with a catalyst. The red line is the catalyzed pathway that has a lower activation energy. This will speed up the reaction.
Diagram of chemical reaction pathway also showing the catalyzed pathway

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