Lesson Explainer: Catalysts Chemistry • 7th Grade

In this explainer, we will learn how to explain the effects of catalysts on the rate of reaction.

Many reactions in industry, and even in the cells of living organisms, would occur too slowly to be useful without something to speed them up. Chemical manufacturers therefore often use catalysts to speed up the rate of reactions, and the cells in living organisms use biological catalysts called enzymes to do the same. A catalyst is a substance that increases the rate of a chemical reaction but that is, itself, not changed chemically by the end of the reaction.

For a chemical reaction to occur, the reactant particles first need to collide with each other. Sometimes, even when particles collide, the collision might not result in a chemical reaction. For example, nitrogen gas and hydrogen gas do not spontaneously react under normal conditions when their particles collide:

and gases do not react under normal conditions because particles need to collide with sufficient energy to react, and this does not happen under normal conditions. They need a minimum amount of energy for a reaction to occur. This minimum amount of energy needed for particles to react with each other when they collide is called the activation energy.

If particles collide with enough energy and are in the correct orientation, bonds in the reactants will break and new bonds will be formed to produce the product ammonia. This is a reversible reaction. The enthalpy change () for the forward reaction is negative, so the forward reaction is exothermic:

The activation energy required for and gases to react is high, so to induce and gases to react, a lot of energy must be supplied. This could be achieved by heating up the reaction system greatly, by pressurizing it, or by reducing the required activation energy by adding a suitable catalyst.

In industry, an iron catalyst is used in the reaction of nitrogen gas and hydrogen gas to produce ammonia. This process is called the Haber–Bosch process.

The activation energy barrier is overcome by the addition of the iron catalyst. Catalysts act by reducing the amount of activation energy needed for a successful collision between the reactant molecules. The reaction profile below shows two activation energy “humps”—one without a catalyst and one with a catalyst, for the forward, exothermic, reaction.

Notice that the activation energy for the forward reaction is lower when the iron catalyst is present: with catalyst < without catalyst.

As a result of the presence of a catalyst, less energy is required for a successful collision between the nitrogen and hydrogen reactant molecules. This causes more successful collisions to occur more frequently, at a given temperature. This is how the rate of the reaction is increased by the presence of a catalyst.

The red line in the diagram shows an “alternative reaction pathway.” In this pathway, the reactant particles are either oriented in a way that increases the likelihood of successful collisions or temporarily bound to the catalyst, forming an intermediate that has a lower energy requirement to form the products than the reactants alone. These mechanisms are two examples of an alternative pathway for a catalyzed reaction and offer different steps to an uncatalyzed reaction.

We can subtract the with a catalyst from the without a catalyst:

The difference, , is activation energy reduction as a result of the catalyst activity.

The difference in energy between the products and the reactants, , is the same for both the catalyzed and uncatalyzed reactions. Catalysts only influence the activation energy required for the reaction to occur but not the enthalpy change.

The reaction profile below is for the reverse reaction. It shows the activation energy with and without the iron catalyst.

This reaction profile shows us that without catalyst > with catalyst is also true for the reverse, endothermic, reaction.

Example 2: Interpreting Reaction Profiles with and without Catalysts

In the following two diagrams, which two letters indicate reaction pathways that have been catalyzed?

Answer

A catalyst acts to speed up the rate of a reaction by lowering the activation energy needed for the reactants to react with each other. In the first reaction profile above, pathway B has a lower activation energy than pathway A. B is the catalyzed reaction, and A is the uncatalyzed reaction.

In the second reaction profile, pathway D has a lower activation energy than pathway C. D is the catalyzed reaction, and C is the uncatalyzed reaction.

The correct answer is B and D.

The diagram below shows how and gas molecules are converted to ammonia by alternative reaction steps as a result of the presence of the iron catalyst.

In step 1, the reactant molecules adsorb onto the iron catalyst surface and dissociate, or break apart, into atoms of nitrogen and atoms of hydrogen. In step 2, three individual hydrogen atoms react with and bond to a nitrogen atom, one at a time, to form an ammonia molecule. In the last step, ammonia is released, or desorbed, from the catalyst surface.

The step-by-step addition of hydrogen atoms to the nitrogen atom is an alternative pathway with lower energy requirements than the reaction of and molecules with each other without a catalyst.

Although a catalyst temporarily bonds to the reactant particles during the reaction, at the end of the reaction, the catalyst returns to its original state. Catalysts do not take part in the reaction and are unchanged chemically by the end. This makes catalysts reusable.

Very small amounts of catalyst are needed because the same catalyst particles are used over and over again. Increasing the amount of catalyst does not necessarily make a significant difference to a reaction, as high efficiency is achieved from just a small amount of catalyst. It is the surface area of the catalyst, or the number of catalyst particles exposed to reactant particles, that can appreciably influence the rate of a reaction. Increasing the surface area of a solid catalyst by grinding it into a fine powder, or working it into a fine mesh, will increase the number of catalyst particles exposed to reactant particles and will therefore increase the rate of the reaction.

Example 3: Understanding How Catalysts Affect the Rates and the Yields of Reactions

Experiment A relies upon a catalyst. If a greater quantity of the catalyst is used, what changes to the graph would be most likely?

1. The graph would be the same shape, and the final volume of gas collected would be lesser.
2. The graph would initially be less steep, and the final volume of gas collected would be the same.
3. The graph would be the same shape, and the final volume of gas collected would be greater.
4. The graph would initially be steeper, and the final volume of gas collected would be the same.
5. The graph would initially be less steep, and the final volume of gas collected would be lesser.

Answer

Only a small amount of catalyst is needed to speed up the rate of a chemical reaction significantly. Adding more catalyst will increase the rate only slightly; thus, in this reaction, adding more catalyst will initially make the graph steeper.

Catalysts do not influence how much product is produced, but only speed up the rate at which the product is made. In this reaction, the volume of gaseous product collected would be the same regardless of whether little or a lot of catalyst is used.

The correct answer is option D, the graph would initially be steeper, and the final volume of gas collected would be the same.

The formula of the catalyst is not usually included in a chemical equation since catalysts are unchanged chemically at the end of a reaction and catalysts are not considered as part of the reactants or products. However, sometimes the catalyst formula is written above the arrows; for example, Fe can be written above the arrows for the iron catalyst in the Haber–Bosch process:

Using catalysts is highly beneficial to industrial manufacturers. Firstly, because catalysts speed up reactions, time is saved. Reactions can be carried out faster when a catalyst is present and more of the desired product can be produced in a given amount of time. In this way, catalysts make the production of substances more efficient. Secondly, reactions that would normally need very high temperatures without a catalyst need less energy input when a catalyst is used. Energy and money are saved because high temperatures are not necessary when a catalyst is present.

Some reactions, however, require very specialized catalysts that are expensive. Besides iron, many of the transition elements are effective catalysts in very small quantities but are costly. Rhodium, palladium, and sometimes platinum are examples of commonly used transition metal catalysts.

In reality, over a period of time, catalysts may become contaminated or “poisoned” by bonding to certain chemical species; for example, iron bonds strongly to sulfur. If sulfur contaminants are present in the Haber–Bosch process, the iron catalyst can become poisoned by bonding with the sulfur compounds. The catalyst then becomes less efficient and will need to be extracted and cleaned or replaced. To prevent contamination of expensive catalysts, manufacturers endeavor to use pure reactant feeds.

The table shows some important catalysts in industry and in our daily lives.

Most of these catalysts are produced from transition metals. One of these examples is the use of platinum, rhodium, and palladium in catalytic converter units in modern vehicles, where the catalysts in the catalytic converters reduce harmful compounds in exhaust gases to safer, less toxic compounds. The picture below shows the catalytic converter unit, which is a part of the exhaust system.

Cells of living organisms catalyze reactions using biological protein catalysts called enzymes.

The diagram below shows how an enzyme has a specific active site to which a substrate (reactant) molecule or molecules can bond. Held in the active site position, the reaction occurs with lower activation energy and the product or products are released from the enzyme, leaving the enzyme in its original state.

An example of an enzyme used by cells to catalyze a reaction is the enzyme catalase. Most living organisms produce this enzyme. Catalase protects the organism from oxidative damage by hydrogen peroxide (), which is a substance produced during metabolic processes. The enzyme converts to less harmful products—water and oxygen.

does decompose on its own, without a catalyst, but very slowly. The catalase enzyme causes the decomposition process to occur much faster, thus reducing the time that cellular components are exposed to oxidizing .