Chemistry
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Catalysts

Catalysts

If you've been studying chemistry, you've probably heard a lot about catalysts. Catalysts are substances that can make chemical reactions happen faster, without getting used up themselves. They work by lowering the amount of energy needed to start the reaction, which is called the activation energy.

Transition metals are one type of catalyst that is especially effective. In this article, we'll explain how they work. We'll also talk about the two main types of catalysts: heterogeneous and homogeneous. You'll learn how these different catalysts are used in industry, including in the Haber and Contact processes. Finally, we'll talk about a special type of reaction called autocatalysis. By understanding how catalysts work, you can gain a better appreciation for the amazing world of chemistry. So let's dive in!

Catalysts

Catalysts are substances that speed up chemical reactions without being consumed themselves. In other words, they remain chemically unchanged and the amount of catalyst produced at the end of the reaction is the same as the amount added at the beginning. This makes them incredibly useful in industry, where they can help to increase the reaction rate and lower the activation energy required.

Transition metals are particularly effective as catalysts due to their variable oxidation states. They can adsorb substances on their surface and activate them, leading to faster reactions.

There are two types of transition metal catalysts: heterogeneous and homogeneous. Heterogeneous catalysts are in a different phase than the reactants, while homogeneous catalysts are in the same phase. Understanding the difference between these types of catalysts can help us understand how they work and how they are used in different industries. To learn more about the factors affecting reaction rates and how catalysts play a role, check out our article on the topic.

Heterogeneous Catalyst

Heterogeneous catalysts are in a different phase than the reactants, and the reaction takes place at active sites on the catalyst's surface. Most heterogeneous catalysts are in the solid phase and do not get consumed in the reaction. Instead, at least one of the reactants gets adsorbed at active sites on the catalyst's surface.

One example of a heterogeneous catalyst in action is the catalytic converter in cars. In this case, a gaseous reactant passes over a solid catalyst, which helps to reduce harmful emissions.

Transition metals are popular solid catalysts, and they are used in the Haber and Contact processes to increase reaction rates. In the Haber process, for example, hydrogen and nitrogen from the air are placed in a reactor with iron pellets as a catalyst. The surface of the iron attracts the hydrogen and nitrogen gas molecules, which become adsorbed onto the surface. The reaction takes place while the molecules are on the surface of the iron, and when ammonia forms, it desorbs from the surface. This process produces ammonia through heterogeneous catalytic action, and it is used in the production of fertilisers around the world.

The Hider process

The Haber process is a reversible process, and the purpose of iron as a catalyst is to speed up the reaction without affecting the equilibrium. Without the iron catalyst, the process would take too long to produce ammonia. This process was invented by German scientist Fritz Haber in 1908, and it revolutionized agriculture. However, only half of the nitrogen in fertilizers is assimilated by plants, leading to high concentrations of ammonia in our water supply and the atmosphere.

Chemists are now searching for alternative processes that can be used to produce abundant crops without endangering the planet.

To understand more about the Haber process, you can learn about compromise pressure in Le Chatelier's Principle. Additionally, adsorption and desorption are essential concepts in heterogeneous catalysis. Adsorption happens when the reactants stick to the catalyst's surface so that the reaction can take place. Desorption happens when the product molecules leave the surface of the catalyst. These processes are essential for understanding how catalysts work and how they can be used in different industries.

The Contact Process

Sulfuric acid is a crucial component in the production of many industrial products such as dyes, detergents, paints, plastics, fertilisers, and fabrics. Globally, we manufacture 231 million tonnes of sulfuric acid every year, with most of it being used to make fertilisers. The Contact process is another example of heterogeneous catalysis that is used to make sulfuric acid industrially.

The Contact process occurs in three stages, with the second stage involving the use of a heterogeneous catalyst, vanadium pentoxide (V2O5), to speed up the reaction rate. The reaction involves the reaction of sulfur dioxide with oxygen to produce sulfur trioxide. The reaction is reversible, and the vanadium pentoxide works to speed up the reaction rate, making the process more efficient.

In the Contact process, sulfur dioxide and oxygen are introduced into the reactor as gases, and they pass over the solid vanadium pentoxide catalyst. Surface adsorption theory explains how the catalyst works, with adsorption happening when one of the reactants attaches to the catalyst's surface. The reaction occurs while the reactants are on the catalyst surface, and desorption happens when the product of the reaction detaches from the catalyst's surface.

The reaction stages of the Contact process illustrate the surface adsorption theory in action. In stage one, sulfur dioxide adsorbs onto the vanadium pentoxide catalyst, and a redox reaction occurs as the catalyst's vanadium is reduced from +5 to +4. The sulfur trioxide desorbs from the catalyst's surface. In stage two, oxygen reacts on the catalyst's surface, causing another redox reaction, which oxidises the vanadium pentoxide back to +5, regenerating the original catalyst, V2O5.

The Contact process
The Contact process

The catalytic converter is one of the most significant advancements in reducing pollution from vehicle exhaust gases. Carbon monoxide and nitrogen monoxide are major pollutants produced by cars, and they pose a risk to both the environment and human health. Catalytic converters work by using a ceramic honeycomb coated with a solid metal catalyst, usually a mixture of platinum and rhodium, to reduce harmful exhaust gases. Surface adsorption theory explains how the transition metals act as catalysts, with harmful gases adsorbing onto active sites where they react to produce harmless gases. The products then desorb and are released through the car exhaust.

The honeycomb interior of the catalytic converter increases its efficiency by providing a greater surface area for the solid catalyst, meaning there are more active sites for adsorption to occur. This structure also helps to minimise the cost of catalytic converters since platinum and rhodium are expensive metals.

However, catalytic converters have a drawback in that they can become "poisoned" by impurities that block active sites. Lead compounds in leaded petrol are one example of a catalyst poison that can "poison" platinum and rhodium, reducing efficiency and increasing the cost of the process. Homogeneous catalysts, on the other hand, are soluble in the same phase as the reactants. These catalysts work by either donating or accepting electrons from the reactants, thereby lowering the activation energy and increasing the reaction rate. Homogeneous catalysts have several advantages over heterogeneous catalysts, such as better selectivity and the ability to control the reaction conditions precisely. However, they also have their drawbacks, such as the difficulty in separating them from the reaction mixture, and the fact that they can be expensive or toxic.

Homogeneous Catalyst

A homogeneous catalyst is a catalyst which is in the same phase as the reactants.

Homogeneous catalysis often involves an aqueous catalyst and aqueous reactants, but this is not always the case. Sometimes, the catalyst and the reactants will be in the gas phase. In homogeneous catalysis, the reaction proceeds through an intermediate species. What does this mean? You can see how this works in the reaction between the persulfate ion (peroxydisulfate) and iodine.How Fe2+ ions catalyse the reaction between iodide and persulfate ionsPersulfate (or peroxodisulfate, its IUPAC name) acts as an oxidising agent when it reacts with iodide. The equation for the reaction is given below:S2O82- + 2I- ➔ 2SO42- + I2At room temperature, the process is slow due to the negatively charged reactant ions repelling each other. However, in the presence of Fe2+ ions, the reaction is much faster. The Fe2+ ions and the reactants are in the aqueous phase, so this reaction is an excellent example of a homogeneous catalyst. Let us look at the steps below:1. Since the Fe2+ ions and the S2O82- ions share opposite charges, they attract each other and react as follows:  S2O82- + 2Fe2+ ➔ 2SO42- + 2Fe3+2. The Fe3+ ions produced in the first reaction react with the I- ions as follows: 2Fe3+ + 2I- ➔ 2Fe2+ + I2As you can see, the original iron(II) ion catalyst gets regenerated, so the steps repeat themselves. The iron(III) ions that form in the first reaction act as the intermediate species.In the process between iodide ions and persulfate ions, we can also use iron(III) ions as the original catalyst. In this case, Fe2+ is the intermediate species. The reaction would take place in the reverse order:

2Fe3+ + 2I- ➔ 2Fe2+ + I2S2O82- + 2Fe2+ ➔ 2SO42- + 2Fe3+

Nitrogen dioxide (NO2) as a catalyst.An example of the use of a homogenous catalyst can be seen with nitrogen dioxide. We can explore this through formation of acid rain which contains H2SO4 (sulphuric acid).When sulphur dioxide (SO2), a pollutant in the atmosphere, is oxidised, it turns into SO3. This can then react with rain water to produce H2SO4. This can be shown as:SO3(g) + H2O(l) → H2SO4(aq)So what role does nitrogen dioxide play in this reaction? It not only leads to acid rain, but also acts as a catalyst. This is due to nitrogen dioxide catalysis of SO2 to SO3.This can be shown as:SO2(g) + NO2(g) → SO3(g) + NO(g)The NO in this reaction can be regenerated, like most catalysts, back to NO2 where it can go on to catalyse the reaction of SO2 to SO3.This can be presented as:NO(g) + 1/2 O2(g) → NO2(g)

Before we conclude this discussion on catalysts, let us look at one last type of catalysis, namely, autocatalysis.

Autocatalysis

In the following reaction, negative manganate(VII) ions react with negative ethanedioate ions.

2MnO4- (aq) + 5C2O42- (aq) + 16H+ (aq) ➔ 2Mn2+ (aq) + 10CO2 (g) + 8H2O (l)

This reaction is fascinating, as the reaction rate increases as more Mn2+ ions get produced. We call it autocatalysis when a process is catalysed by one of the products of the reaction. In this case, Mn2+(aq) acts as an autocatalyst. The process starts slow, but as more manganese (II) ions form, it becomes faster and faster. Eventually, the reaction slows down when the catalyst gets used up.

Similar to the previous reaction where iron acts as a catalyst, the Mn(II) gets regenerated in a redox cycle.

The Mn2+ ions react with the MnO4- ions to produce Mn3+ ions.The Mn3+ ions react with the ethanedioate ions to regenerate Mn2+ ions.

We can show this redox cycle using the following equations:

4Mn2+(aq)  +  MnO4-(aq) + 8H+(aq)   →   5Mn3+(aq)  + 4H2O (aq)

2Mn3+(aq)  +  C2O42-(aq)  →  2CO2(g) +  2Mn2+(aq)

And there you have it: the wonderful world of transition metal catalysts!

 

To summarise, a catalyst is a substance that speeds up a chemical reaction without being changed in composition or quantity. There are two types of catalysts: heterogeneous catalysts and homogeneous catalysts.

Heterogeneous catalysts are in a different phase from the reactants, and they work through surface adsorption theory. The Haber process and the Contact process both use heterogeneous catalysts to produce ammonia and sulfuric acid, respectively. Catalytic converters also use heterogeneous catalysts to reduce vehicle emissions.

Catalytic converters are typically ceramic honeycombs coated with a mixture of platinum and rhodium or palladium, which act as catalysts through surface adsorption theory. However, catalytic poisoning can reduce their efficiency.

Homogeneous catalysts are in the same phase as the reactants, and they work by either donating or accepting electrons from the reactants, thereby lowering the activation energy and increasing the reaction rate. Homogeneous catalysis has several advantages over heterogeneous catalysis, such as better selectivity and the ability to control reaction conditions precisely.

Finally, autocatalysis occurs when a process is catalysed by one of the reaction products.

Catalysts

What are the 3 types of catalysis?

The three types of catalysis are:Homogeneous catalysisHeterogeneous catalysisAutocatalysis 

How do you identify a homogeneous catalyst?

A homogeneous catalyst is in the same phase as the reactants. Homogeneous catalysis often involves an aqueous catalyst and aqueous reactants, but this is not always the case. Sometimes, the catalyst and the reactants will be in the gas phase. In homogeneous catalysis, the reaction proceeds through an intermediate species.

How do you identify a heterogeneous catalyst?

A heterogeneous catalyst is in a different phase from the reactants, and the reaction occurs at active sites on the surface.In general, most heterogeneous catalysts are in the solid phase and do not get consumed in the reaction. At least one of the reactants gets adsorbed at active sites on the catalyst's surface in heterogeneous catalysis.

How do catalysts speed up reactions?

Catalysts speed up a reaction by lowering the required activation energy. 

What is a catalyst?

A catalyst is a substance that increases the rate of a chemical reaction without being changed in chemical composition or amount.

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