Catalyst - Online Tutor, Practice Problems & Exam Prep

A catalyst is a substance that accelerates a chemical reaction by lowering the activation energy without being consumed by the reaction itself. The energy diagram illustrates this by showing a lower energy barrier for the catalyzed reaction compared to the uncatalyzed one, indicating a reduced activation energy. The height of the activation energy is inversely related to the reaction speed; a lower activation energy means reactants can more readily overcome the energy barrier and transform into products. This concept is crucial in understanding how catalysts influence the dynamics of chemical reactions, making them faster and more efficient.

Catalyst Concept 1

A catalyst can be seen as any substance that increases the rate of reaction by decreasing your energy of activation and also not being consumed in the process. Here we have a catalyzed versus an uncatalyzed reaction being displayed in terms of an energy diagram. Remember, in our energy diagram we have our reactants in the beginning, our products at the end. The difference between the two lines represents the overall change in energy of the chemical reaction.

And then if we take a look here, we have at the top of this first hill is our transition state for it. And at the top of this hill is our transition state. The difference between a transition States and the reactant line gives us our activation energy. Remember, a catalyst lowers your energy of activation. If we take a look at the first red curve, we can see that its activation energy is pretty high. So it's safe to assume that this is the energy of activation for our uncatalyzed reaction. So here we'd say that this is EA uncapped.

Now when we add a catalyst, the energy of activation is expected to go down. So the line in blue represents the addition of our catalyst. So this would be EA cap per catalyzed. So remember the height of the activation energy has a direct impact on the speed of the reaction. The higher the activation energy, the longer it takes for my reactants to basically climb and scale this hill to slide down and become a product.

So the red line here we say since it's higher up and higher activation energy, it would take longer for us to make our products the catalyzed one. Since the activation energy is smaller, we expect the reactants to have an easier time getting up this hill and sliding down to make our products. So keep in mind this is what a catalyst does. It lowers our energy of activation so that our reactants can more easily become products.

Catalyst Example 1

Hey, everyone. So in this example question, it says a certain reaction has an enthalpy value of -20 kilojoules and an activation energy of 40 kilojoules. A catalyst is found that lowers the activation energy of the reaction by 10 kilojoules. What is the total difference in energy between the products and the transition state?

All right, so they're telling us here our enthalpy value. So this is ΔH and we can say here ΔE also can fill in and it's -20. So let's say that we're our reactants start here at 40 kilojoules, and -20 means our products are 20 kilojoules lower, so they're here at 20 kilojoules. Activation energy. Remember, activation energy is the difference in height between my transition state and my reactant line. So that would mean my transition state which is up here would have to be 80. The difference between these two is what my activation energy is.

Now remember ΔH, Oregon ΔE equals products minus reactants. So it D be 20, -, 40, which is how we got that -20, all right. So now they're saying that we used a catalyst which lowers the energy of activation. So here we were at 80 for the transition state, it gets lowered by 10. So now we're at 70. We're still at 40 for the reactants. That doesn't change and my products are still down here at line 20.

So now we have to calculate the difference. This ΔH or ΔE doesn't change. It's still going to be -20 and then here differences in energy between products and the transition state ear if we're looking at what's the total difference in energy between products and transition state. So we're looking at this total distance here, right. So here we'd say total difference in energy between products and the transition state, products are at 20, transition state is 70. So that would be a difference of 50 kilojoules total.

OK. So that's the total difference. We start out with a total difference of 60 kilojoules between products and transition state. We add the catalyst, the overall energy is lowered by that much. It'd be 50 here. If we wanted to look at the new activation energy, we could also do that. The activation energy here, the new one, would be 70 minus the 40, so it would be 30 kilojoules. And again ΔH or ΔE would still be -20 kilojoules. So this is all the information we could tell for this particular energy diagram once we've introduced the catalyst.

Which of the following statements is true regarding the energy diagram provided? 

i. The reaction is endothermic.

ii. Activation energy would be less than + 10 kJ after a catalyst is added.

iii. The reaction absorbs energy.

iv. Activation energy would be greater than + 10 kJ after a catalyst is added.

• A. I only
• B. II only
• C. I and III
• D. II, III and IV

Here’s what students ask on this topic:

What is a catalyst?

A catalyst is a substance that increases the rate of a chemical reaction without itself undergoing any permanent chemical change. It works by providing an alternative reaction pathway with a lower activation energy compared to the uncatalyzed mechanism. This means that more reactant particles have enough energy to react when a catalyst is present, leading to an increased rate of reaction.

Catalysts are crucial in various industrial and biological processes. In industry, they are used to enhance the efficiency of chemical production, such as in the synthesis of fertilizers, plastics, and fuels. In biological systems, enzymes act as natural catalysts that facilitate essential biochemical reactions necessary for life, such as digestion and metabolism.

Catalysts can be heterogeneous, where they are in a different phase than the reactants, or homogeneous, where they are in the same phase as the reactants. They are also characterized by their ability to be recovered and reused multiple times without losing their effectiveness.

How does a catalyst influence a chemical reaction?

A catalyst is a substance that increases the rate of a chemical reaction without being consumed in the process. It works by providing an alternative reaction pathway with a lower activation energy compared to the uncatalyzed mechanism. Activation energy is the minimum energy required for reactants to undergo a chemical reaction. By lowering this energy barrier, a catalyst allows more reactant particles to have enough energy to react when they collide, thus speeding up the reaction.

Catalysts do not alter the equilibrium position of a reaction; they only make it faster to reach. They are specific in nature, meaning that a particular catalyst will only work for a specific reaction or set of reactions. Catalysts are widely used in industry to increase the efficiency of chemical processes, leading to cost savings and lower energy consumption. In biological systems, enzymes act as natural catalysts to facilitate complex biochemical reactions under mild conditions, which are essential for life processes.

Which component is affected when a catalyst is added to a chemical reaction?

When a catalyst is added to a chemical reaction, it primarily affects the reaction rate. A catalyst speeds up the rate at which the reaction proceeds without being consumed in the process. It does this by providing an alternative reaction pathway with a lower activation energy than the uncatalyzed mechanism. This means that more reactant particles have the necessary energy to undergo the transformation into products at any given time, increasing the frequency of successful collisions. The catalyst remains chemically unchanged after the reaction, so it can continue to facilitate additional reactions. It's important to note that while a catalyst accelerates the reaction rate, it does not alter the equilibrium position of a reversible reaction; it simply helps the system reach equilibrium faster.