In this video, we will learn what a catalyst does in terms of its effect on the rate of a chemical reaction. We will also learn about how a catalyst works and the reasons why catalysts are used in industrial chemical processes.
Chemical reactions usually happen when reactor molecules are mixed together and collide with each other. For example, we could mix methane gas with oxygen gas. The methane and oxygen gas particles move around freely, and the molecules will collide with each other frequently. But a chemical reaction does not occur under normal conditions. In this case, we do not automatically get combustion or burning taking place. We need to add a spark or a flame to get the reaction started with a bit of a bang.
On the other hand, some chemical reactions seem to start all by themselves. White phosphorus is a dangerous element to leave lying around in the presence of air. It’s normally stored under water. As it dries out in air, it will first smolder and then ignite all by itself. It automatically reacts with oxygen in the air to produce phosphorus oxide. In order to understand what makes chemical reactions happen, we clearly need to explore the reasons behind why some reactions start easily all by themselves, whereas others need heating, sometimes to high temperatures, first.
To get a chemical reaction to happen, we need reactant particles to collide with each other first. If we mix methane gas and oxygen gas together, there’s no shortage of collisions between these molecules every second because they’re both in the gas state. But we also need these particles to collide with just enough energy to break chemical bonds within these reactant molecules. In these collisions, the minimum energy needed to break bonds and cause what we call a successful collision is known as the activation energy. Activation energy is the minimum energy needed for a reaction to occur.
In many reactions, including the reaction of methane gas with oxygen gas, this activation energy is very high. So at normal temperatures, reactive particles just bounce off each other when they collide. Bonds are not broken, and a chemical reaction does not take place. Most collisions at normal temperatures are not successful collisions. In the reaction of methane and oxygen, the chemical energy stored in the reactants is greater than the chemical energy stored in the products.
The difference between the energy level of the reactants and the energy level of the products is called the energy change for this reaction. The energy change is often seen labeled as ΔH. This is a large triangle symbol with a capital H. The energy change for the reaction of methane with oxygen is negative as heat is released to the surroundings. A reaction where heat is released to the surroundings is described as an exothermic reaction. All combustion reactions are exothermic. The reaction of methane with oxygen is a combustion reaction.
In the reaction of methane with oxygen, which is an exothermic reaction releasing heat to the surroundings, the activation energy is high enough to ensure that methane and oxygen cannot automatically or spontaneously react with each other at normal temperatures. We need to put some additional energy into the reactant mixture first to provide the activation energy to break bonds and get the reaction started. This additional quantity of energy is labeled as the activation energy on our energy level diagram here. The reaction then proceeds automatically or spontaneously as it releases enough heat energy in the process to maintain successful collisions.
The situation is similar for an endothermic reaction, where heat is absorbed from the surroundings and the products finish at a higher energy level relative to the reactants that we started with. Endothermic reactions, like all chemical reactions, also require activation energy to get them started. We must raise the energy level of the reactants above that of the products to provide this activation energy. This is labeled on the energy level diagram.
So in summary at this point, some reactions are harder to get started than others because they have a high activation energy. Some reactions happen very slowly or not at all under normal conditions because they have a high activation energy. How can we make these reactions with a high activation energy happen at all?
Hydrogen peroxide is a liquid substance that can decompose to make water and oxygen gas. This reaction is very slow under normal conditions. You wouldn’t see many bubbles of oxygen gas forming in the liquid hydrogen peroxide. If we add a small amount of a black powder called manganese dioxide to the hydrogen peroxide, we see a sudden increase in the number of oxygen bubbles being produced every second. The manganese dioxide is increasing the rate at which bubbles of oxygen are produced. It is increasing the rate of reaction very dramatically in this case.
We have gone from a situation where a few bubbles of oxygen gas are being produced every now and then to many bubbles of oxygen gas being produced in a very short space of time. The manganese dioxide is behaving as a catalyst. A catalyst is a substance that increases the rate of a reaction without itself undergoing a permanent chemical change. The energy profile diagram for the decomposition of hydrogen peroxide shows us that the overall energy change is negative. The decomposition of hydrogen peroxide is therefore an exothermic reaction. This is the case if the reaction goes ahead with or without a catalyst.
As with all chemical reactions, there is an activation energy associated with this decomposition process. If we add a catalyst to the reactant, we find that the activation energy is much smaller than it is without a catalyst. Catalysts speed up the rate of a chemical reaction. They do this by offering an alternative reaction pathway or route for the reaction to take place. Most importantly, this alternative reaction pathway has a reduced or lower activation energy. With a lower activation energy, a much greater percentage of the collisions involving reactant particles will lead to bonds breaking and a reaction occurring. More successful collisions occur every second, and the rate of reaction will be increased.
This situation is similar to going on a journey over Mont Blanc, the highest mountain in France. We could expend a lot of energy climbing over the top of Mont Blanc and visiting Italy on the other side. This is the high energy route. Alternatively, we could pass through the tunnel quite easily instead. This would be a lower energy route and would be much faster.
We now know that catalysts speed up chemical reactions by providing an alternative reaction pathway with a lower activation energy. So how will the catalysts become involved in a chemical reaction? Although catalysts remain chemically unchanged at the end of the reaction that they are catalyzing and they can be recovered, it will be incorrect to say that catalysts do not become chemically involved or changed during the reaction that they are catalyzing.
Catalysts can concentrate reactant particles on their surface. These reactant particles would be moving around randomly and occasionally colliding with each other in the process. By doing this, catalysts weaken bonds in the reactant particles. This is how the activation energy for the reaction is reduced. At the end of the reaction, products are released and the catalyst is chemically unchanged. We could say that it is chemically regenerated. It’s not used up or consumed as a reactant normally is during the chemical reaction. For this reason, the overall equation for the chemical reaction will not usually include the catalyst. The catalyst may be written above the arrow instead.
As catalysts are not consumed during a chemical reaction, only a small quantity of catalyst need be used for a given situation. In a catalytic converter found within a car exhaust pipe close to the engine, toxic gases such as carbon monoxide and oxides of nitrogen enter the catalytic converter. These toxic gases are chemically changed to safer gases such as carbon dioxide and nitrogen, which exit the tailpipe.
In many situations, such as the catalytic converter in car exhaust pipes, a small amount of catalyst is spread over a large surface area. At normal operating temperatures, it will be very hot. It’s important to remember here that raising the temperature of a reaction increases the rate much more than using a larger quantity of catalyst. Both of these factors, surface area and temperature, help to improve the efficiency of the catalyst whilst keeping the amount used to a minimum.
Catalysts often contain transition metals or elements located in the central block of the periodic table. Some of these metals are highly expensive, such as platinum and palladium, which are used in the catalytic converter. These metals can, of course, be recovered at the end of the car’s life. Although some common industrial catalysts contain precious metals, many catalysts contain more common metals and they’re relatively cheap.
In the Haber process, ammonia is produced. Ammonia is an important starting material for fertilizers, which help to improve crop yields. In the Haber process, nitrogen gas is reacted with hydrogen gas. Nitrogen and hydrogen do not normally react with one another unless high temperature and very high pressures are used. In the Haber process, iron is used as a catalyst to lower the activation energy and allow the reaction to proceed at lower temperatures with a reasonable rate of reaction.
The situation is complicated by the fact that this reaction is reversible. The iron catalyst speeds up the forward and reverse reaction rates equally, allowing the mixture to reach equilibrium in the reactor vessel in less time. Some ammonia is obtained in less time, regardless of the yield for this reversible process. Any ammonia formed in the reactor vessel is cooled, liquefied, and collected. Unreacted nitrogen and hydrogen are recycled and returned to the reactor vessel.
Catalysts save vast amounts of energy in industry by allowing reactions to proceed at lower temperatures with a reasonable rate. This saves money, in terms of the expensive energy saved, and time, since the reaction will be taking place at a faster rate than it would without a catalyst. The economic benefit of using a catalyst is substantial. And it will often outweigh the cost of the catalyst.
Now let us look at a question to test our understanding of catalysts.
Manganese dioxide is used as a catalyst in the decomposition of hydrogen peroxide to form water and oxygen. Which statement is untrue when using MnO2 as a catalyst? 2H2O2 aqueous produces 2H2O liquid plus O2 gas. (A) The oxygen will be formed more quickly. (B) More oxygen will be produced. (C) The mass of MnO2 before and after the reaction will be the same. (D) An alternative reaction pathway is provided by the catalyst. (E) The catalyst remains unchanged at the end of the experiment.
In this question, we’re being asked about how a catalyst, which is manganese dioxide, behaves in a decomposition reaction. In this decomposition reaction, a single reactant, the hydrogen peroxide, decomposes to form new products. The products are water and oxygen gas. This decomposition reaction is very slow under normal conditions at room temperature. We would not observe many bubbles of oxygen gas coming from the hydrogen peroxide solution at all. This situation would change rapidly if a small amount of the solid catalyst, manganese dioxide, were added to the hydrogen peroxide solution.
Rapid fizzing or effervescence would be observed as soon as the black powder is added to the hydrogen peroxide. The manganese dioxide catalyst will increase the rate of decomposition of the hydrogen peroxide. More oxygen gas, seen as bubbles, will be produced per unit of time as the reaction rate has been increased. Oxygen gas will certainly be produced more rapidly. This statement is true, so it’s not the correct answer. Remember, in this question, we’re looking for an untrue statement.
Notice that the oxygen gas, which is one of the products, originates from the hydrogen peroxide molecules. According to the balanced equation, two molecules of hydrogen peroxide are required to produce one molecule of oxygen gas. If we have a fixed amount of hydrogen peroxide molecules at the start of the reaction, we can only produce a fixed amount of oxygen molecules during the decomposition reaction. Adding the manganese dioxide catalyst does not change the amount of oxygen gas produced. It simply increases the rate of reaction.
The same amount of oxygen gas is produced in much less time. The manganese dioxide does not appear in the overall reaction equation. The amount of oxygen gas obtained will be the same with or without the catalyst present. This means that the statement that more oxygen gas is produced is untrue. Therefore, it’s likely to be the correct answer.
Let us consider the other possible responses first before committing to this response. We could take one gram of manganese dioxide catalyst, record its mass, and add it to the hydrogen peroxide at the start of the reaction. When the reaction is complete, the beaker will contain only pure liquid water and manganese dioxide as a mixture. The manganese dioxide could be filtered off, dried, and its mass recorded on an accurate balance. We would find that it is chemically unchanged and, in fact, the mass recovered will be exactly the same as the mass used at the start.
The catalyst MnO2 will be unchanged at the end of the experiment, and the mass of MnO2 before and after the reaction will be the same. Remember that a catalyst becomes involved with a chemical reaction. It may be chemically changed during this process, but it is regenerated at the end of the reaction. Both of these responses are true. And they are therefore not correct answers.
The decomposition of hydrogen peroxide is an exothermic reaction. We can view this in an energy level diagram. And the energy change for this reaction is negative. The MnO2 catalyst increases the rate of the chemical reaction by providing an alternative reaction pathway with a lower activation energy. By reducing the activation energy for the reaction, a greater proportion of reactant molecular collisions are successful. Successful collisions are those where chemical bonds are broken and they lead to a reaction occurring. Hence, we see more successful collisions per second and a faster rate of reaction. The final statement that the catalyst provides an alternative reaction pathway is true and, therefore, not the correct answer here. We’re looking for a statement that is untrue.
The statement that more oxygen will be produced is, therefore, the correct answer to this question.
To summarize, let us review the key points from this lesson. Catalysts increase the rate at which products are formed in a chemical reaction. Catalysts increase the rate of reaction by offering an alternative reaction pathway with a lower activation energy. Catalysts are not consumed in the reaction. And catalysts save time and money in industrial reactions.