Catalysis: Definition, Types of Catalyst and Example

A catalyst is a substance that alters the rate of a chemical reaction, and this process is known as catalysis. The catalyst is not consumed during the catalysed reaction, but it might perform repeatedly. 

The term catalysis was coined by the great Swedish chemist Jons Jacob Berzelius in 1835. He identified certain substances that loosened the bond in the reacting molecules and increased the reaction rate. 

Catalysis is interesting because it exposes the fundamental nature of chemical reactions; in practice, catalysis is essential since many industrial processes rely on catalysts to work. In this article, we will read more about types of catalysts, how they work with different examples. 

What is Catalysis?

In chemistry, catalysis is defined as modifying the rate of a chemical reaction by adding a substance (catalyst) not consumed during the reaction.

A catalyst accelerates a chemical reaction by forming bonds with the reacting molecules and by allowing these to react and form a product, which detaches from the catalyst. It leaves it unaltered such that it is available for the next reaction.

Later, it was identified that many substances even retarded the speed of a reaction. Hence a catalyst is defined as a substance that alters the rate of a chemical reaction without itself undergoing chemical change. The phenomenon which involves the action of a catalyst is called catalysis.

Catalysis is divided into Positive and negative catalysis: In positive catalysis, the reaction rate is increased by the presence of a catalyst, but in negative catalysts, the rate of reaction is decreased by the presence of a catalyst. The two main types of catalysis are
(i) Homogeneous catalysis and
(ii) Heterogeneous catalysis.

How Does a Catalyst Work?

In a chemical reaction, bond making and bond breaking occur. We can correlate this phenomenon with a “square dance where we swap partners”! Sometimes it is easy to break the bonds and form new bonds, but sometimes in very stable compounds, it is quite tough. Hence, in such conditions, we need a catalyst that helps break and rebuild such strong bonds that increase the reaction rate by lowering the activation energy of that particular reaction.

The activation energy (Ea) is the minimum energy required for a chemical reaction to take place. A catalyst lowers the activation energy, but it does not change the energies of the original reactants or products and does not change the equilibrium of a chemical reaction.

Types of Catalysis

Based on the physical state and nature of the substance used in a chemical reaction, catalysis is of two types:

1. Homogeneous catalysis
2. Heterogeneous catalysis

Homogeneous Catalysis

In homogeneously catalysed reactions, the reactants, products and catalysts are present in the same state of matter—for example, Hydrolysis of cane sugar with the mineral acid (sulphuric acid) as a catalyst.

In this reaction, Reactants, products and catalysts all are in an aqueous state. Hence this reaction is an example of homogeneous catalysis.

Heterogeneous catalysis

In heterogeneous catalysed reactions, the catalyst is present in a different phase, i.e. it is not present in the same phase as that of reactants or products. This is generally referred to as contact catalysis, and the catalyst present is in the form of finely divided metal or as gauze.

For example, in the Haber-Bosch process, the reaction involves gas phase reactants and products and a solid catalyst (Iron). The reaction takes place on the surface of the catalyst.

Homogeneous Catalysis vs Heterogeneous Catalysis

Homogeneous catalysis	Heterogeneous catalysis
1. When the reactants and the catalyst are in the same phase (i.e., liquid or gas), the process is said to be homogeneous catalysis. For example, oxidation of sulphur dioxide into sulphur trioxide with di-oxygen in the presence of nitrogen oxides as the catalyst in the lead chamber process.	1. The catalytic process in which the reactant and the catalyst are in different phases and are known as heterogeneous. For example, Oxidation of sulphur dioxide into sulphur trioxide in the presence of Pt.
2. The reaction is assumed to proceed through the formation of an intermediate between the catalyst and one of the reactants.	2. The reaction proceeds through the adsorption of one or more reactants over the surface of the catalyst.

Auto-Catalysis

In some reactions, one of the products formed acts as a catalyst. Initially, the rate of such reaction will be very slow, but with the increase in time, the rate of reaction increases. Autocatalysis is observed in the given reaction where acetic acid acts as the autocatalyst.

Promoters and Catalyst Poison

The promoter will speed up the catalytic activity. It’s a catalyst for a catalyst. In a catalysed reaction the presence of a certain substance increases the activity of a catalyst. Such a substance is called a promoter.
For example in Haber’s process of manufacturing ammonia, the activity of the iron catalyst is increased by the presence of molybdenum. Hence molybdenum is called a promoter.

On the other hand, certain substances when added to a catalysed reaction decrease or completely destroy the activity of the catalyst and they are often known as catalytic poisons.

For example in the Haber’s process of manufacture of ammonia, the activity of the iron catalyst is poisoned in the presence of \({{\text{H}}_{\text{2}}}{\text{S}}\).

Theories of Catalysis

The process of catalysis is explained with the help of two famous theories. Let’s discuss them in detail:

Intermediate Compound Formation Theory

A catalyst acts by providing a new path with low energy of activation. According to this theory, the desired reaction is brought about by a path involving the formation of an unstable intermediate compound, followed by its decomposition into the desired end products with the regeneration of the catalyst. Consider the reactions:

Activation energies for the reactions \(\left( 2 \right)\) and \(\left( 3 \right)\) are lowered compared to that of \(\left( 1 \right).\) Hence the formation and decomposition of the intermediate accelerate the rate of the reaction.

Limitations of Intermediate Compound Formation Theory

(i) The intermediate compound theory fails to explain the action of catalytic poison and activators (promoters).

(ii) This theory is unable to explain the mechanism of heterogeneous catalysed reactions.

Adsorption Theory

The action of catalyst was explained by Langmuir in heterogeneous catalysed reactions based on adsorption. The reactant molecules are adsorbed on the catalyst surfaces, so this can also be called contact catalysis.

According to this theory, the reactants are adsorbed on the catalyst surface to form an activated complex which subsequently decomposes and gives the product. This can be explained with the help of the figure given below:

Enzyme Catalysis

Enzymes are complex protein molecules. They catalyse the chemical reaction in living organisms; hence are also known as biological catalysts. Enzymes are often present in colloidal states and extremely specific in catalytic action.

Each enzyme produced in a particular living cell can catalyse a particular reaction. The place where these substrate molecules fit is called the active site. For example, the Inversion of cane sugar is carried out with the enzyme invertase that converts cane sugar into glucose and fructose.

A diagrammatic explanation of enzyme catalysed reaction is given below:

Summary

Hence, it can be concluded that catalyst increases the reaction rate without getting consumed in the process. Catalysts typically speed up a reaction by reducing the activation energy or changing the reaction mechanism. 

This catalysis work is very similar to a bypass built over a bumpy road that would have a longer route. Enzymes are proteins that act as catalysts in biochemical reactions. In this article, we discussed the types and functions of catalysts.

FAQs

Q.1. What is the theory of catalysis?
Ans. There are two main theories to explain catalysis:

Q.2. What is catalysis explained with examples?
Ans: The process in which the rate of a chemical reaction increases due to the presence of a catalyst is known as catalysis. For example, in Haber’s process of ammonia formation, the reaction involves gas phase reactants and products and a solid catalyst (Iron).

Q.3. What is a catalyst?
Ans: A catalyst is a substance that increases a reaction rate of a reaction without being consumed on its own. In other words, it’s a substance that increases the rate of a chemical reaction without being changed in identity.

Q.4. What is catalysis and its types?
Ans: A catalyst is a chemical substance that accelerates a chemical reaction. The phenomenon which involves the action of a catalyst is called catalysis. Based on this concept catalysis are divided into two main types: (i) Homogeneous catalysis and (ii) Heterogeneous catalysis.

Q.6. What is homogeneous catalysis for example?
Ans: Homogeneous catalysis refers to catalytic reactions where the catalyst is in the same phase as the reactants and products. For example in Haber’s process of preparation of Ammonia, the reactants, products as well as catalysts all are in the gaseous phase.