Catalyst

Catalyst is a common word that you might come across while studying chemistry especially while learning about chemical reactions. While some of the chemical reactions occur quickly, some take a long time and require extra materials or effort. This is where a catalyst comes in. What is a catalyst? We will find out in this lesson.

Table of Content

Catalyst Meaning

In Chemistry, catalysts are defined as those substances which alter the rate of reaction by changing the path of reaction. Most of the time a catalyst is used to speed up or increase the rate of the reaction. However, if we go into a more deeper level, catalysts are used to break or rebuild the chemical bonds between the atoms which are present in the molecules of different elements or compounds. In essence, catalysts encourage molecules to react and make the whole reaction process easier and efficient.

Some of the important characteristic features of catalysts are,

  • A catalyst does not initiate a chemical reaction.
  • A catalyst does not be consumed in the reaction.
  • Catalysts tend to react with reactants to form intermediates and at the same time facilitate the production of the final reaction product. After the whole process, a catalyst can regenerate.

A catalyst can be either solid, liquid or gaseous catalysts. Some of the solid catalysts include metals or their oxides, including sulfides, and halides. Semi-metallic elements such as boron, aluminium, and silicon are also used as catalysts. Likewise, liquid and gaseous elements which are in pure form are used as catalysts. Sometimes, these elements are also used along with suitable solvents or carriers.

The reaction which involves a catalyst in their system are known as catalytic reactions.

Also Read: Chemical Kinetics

Types of Catalysts With Examples

There are several types of catalysts that can be used depending on the need or requirement of the chemical reaction. They are as follows;

Positive Catalysts

Catalysts which increase the rate of a chemical reaction are positive catalysts. It increases the rate of reaction by lowering the activation energy barriers such that a large number of reaction molecules are converted into products, thereby the percentage of yield of products increases.

Example: In the preparation of NH3 by Haber’s process Iron oxide acts as a positive catalyst and increases the yield of ammonia in spite of less reaction of Nitrogen.

Negative Catalysts

Catalysts which decrease the rate of reaction and negative catalyst. It decreases the rate of reaction by increasing the activation energy barrier which decreases the number of reactant molecules to transform into products and hence the rate of reaction decreases.

Example: Decomposition of Hydrogen peroxide into water and oxygen is retarded by using Acetanilide, this acts as a negative catalyst to decrease the rate of decomposition of hydrogen peroxide.

Promoter or Accelerators

A substance which increases the catalyst activity are known as Promoters or accelerators.

Example: In Haber’s process molybdenum or a mixture of potassium and Aluminium oxides act as Promoters.

Catalyst Poisons or Inhibitors

Substances which decrease the catalyst activity are known as catalyst poisons or inhibitors.

Example: In the hydrogenation of alkyne to an alkene, catalyst palladium is poisoned with barium sulphate in quinolone solution and the reaction is stopped at alkene level. The catalyst is known as Lindler’s catalyst.

Also Read: Enzyme Catalyst

Catalysis

When a catalyst is used to increase the rate of a chemical reaction this phenomenon is known as catalysis.

What are the Types of Catalysis?

On the basis of nature and the physical state of substance employed in the chemical reaction, catalysis is of three types;

  • Homogeneous catalysis
  • Heterogeneous catalysis
  • Autocatalysis

Heterogeneous Catalysis

In this type of catalysis, the reacting substances in a reaction and catalyst employed in that reaction are not in the same state of matter.

Examples 1: Preparation of Ammonia by Haber’s process.

Pure and dry Nitrogen and Hydrogen gases in 1 : 3 ratio are passed through a compressor where high pressure of 200 – 30 atmosphere is maintained. In this process, Iron oxide is used as a catalyst. It is solid oxide employed in a process where the reactants are in a gaseous state. The Nitrogen (g) reacts with hydrogen (g) to form Ammonia (g), in the pressure of Iron oxide solid thus it is heterogeneous catalysis

Example 2: Manufacture of sulphuric acid by Contact process.

In this process, oxidation of sulphur dioxide is a major step. In this oxidation, sulphur dioxide is a gas and oxygen is a gas while vanadium pentoxide is a solid catalyst. In this process, reactants and catalyst are in different states of matter.

Mechanism of Heterogeneous Catalyst

Heterogeneous catalysis involves both adsorption as well as intermediate compound formation. Reactant molecule gets adsorbed on the activation centre of the surface of the catalyst. These combine to form an activated complex which is an intermediate compound. This compound decomposes to give products.

As soon as the products formed these get desorbed from the surface without any lapse in time. The heterogeneous catalysis involves initially adsorption of reactants on the surface of catalyst, Intermediate compound formation, dissociating into a product.

Example: Hydrogenation of ethene into ethane in the surface of the nickel.

  • Ether and hydrogen molecules are adsorbed on the surface of the catalyst.
  • Hydrogen occupies most of the activation centre and is known as occlusion.
  • Ethane molecule attack at its double bond region to form an activated complex.
  • Ether reacts with active hydrogen to form ethane.
  • This ethane gets desorbed on the surface of the catalyst.

Hydrogenation of ethene

Homogeneous Catalysis

The catalysis in which the catalyst employed in the reaction and the reactants are in the same state of matter, that process is referred to as homogeneous catalysis.

Example 1: Hydrolysis of ethyl acetate in the presence of dilute acid.

Ethyl acetate is a liquid which contains an ester functional group. It reacts with water in the presence of dilute sulphuric acid which is a liquid to give ethyl alcohol and acetic acid.

\(C{{H}_{3}}COO{{C}_{2}}{{H}_{5}}+{{H}_{2}}OCl \overset{H\oplus Cl} \longrightarrow C{{H}_{3}}COOHCl+{{C}_{2}}{{H}_{5}}OHCl\)

In the above reaction reactants and catalyst are in the same state of matter. Hence, it is homogeneous catalysis

Example 2: Oxidation of sulphur dioxide in the lead chamber process.

Lead chamber process is used in the manufacture of sulphuric acid. In this process, Nitric oxide gas is used as catalysis.

\(2S{{O}_{2}}\left( g \right)+{{O}_{2}}\left( g\right) \overset{NO(g)} \longrightarrow 2S{{O}_{3}}\left( g \right)\)

In the above reaction, SO2 and O2 along with catalyst NO is also a gas hence it is homogeneous catalysis.

Also Read: Wilkinson’s Catalyst

Mechanism of Homogeneous Catalysis

The homogeneous catalysis takes place by intermediate compound formatter theory.

Let us consider the oxidation of SO2 into SO3 by the lead chamber process. In this nitric oxide gas is the catalyst.

This NO reacts with SO2 to form SO2 and “NO2” as an intermediate compound.

\(2S{{O}_{2}}\left( g \right)+{{O}_{2}}\left( g \right) \overset{NO\left( g \right)} \longrightarrow 2S{{O}_{3}}\left( g \right)\)

First step: Nitric oxide combines with oxygen to form nitrogen dioxide (NO2). This NO2 acts as an intermediate compound, which reacts with SO2 to form sulphur trioxide and NO

2NO(g) + O2(g) → 2NO2(g) Intermediate compound

2SO2 + 2NO2 → 2SO3(g) + 2NO(g)

Autocatalysis

In the autocatalytic reaction, there is no specific catalyst that is added. Instead, one of the products acts as a catalyst and increases the rate of formation of products.

Example 1: Decomposition of Arsene (AsH3) is formed by the Arsenic formed in the reactor is “autocatalyst”.

2As H3 → 2As + 3H2

In this process As acts as a catalyst.

Example 2: Permanganate vs oxalic acid during the process Mn+2 formed acts as an autocatalyst.

2mno4 + 5CH2C2O4→2Mn+2+8H2O+16HP

Frequently Asked Questions On Catalyst

How can a positive catalyst alter the reaction?

A positive catalyst is to make reaction rate very first by changing the path of reaction by decreasing the activation energy basis. Such that a large number of reactant molecular converted into products.

What is the role of catalyst poison in Rosenmund reaction?

In Rosenmund reaction Aldehyde are prepared by reducing Acid halides with hydrogen gas in the presence of palladium. If a catalyst is not poisoned the reaction is not stopped at aldehyde level which is feather reduces of alcohol. In order to stop at the aldehyde level. Palladium is poisoned with Barium sulphate.

What are the key factors in heterogeneous catalysis?

In heterogeneous catalysis, the reacting and catalyst are in different states of the matter. The most important steps in this process are;
– Adsorption of reactant molecules activation centre.
– Formation of activation complex at the centre.
– This complex decomposes to give products.
– Desorption of products from the surface of the catalyst.

What is the role of promoters in Haber’s process?

Promotors or accelerators increase the catalyst activity in a process. In Haber’s process of manufacture of Ammonia, Nitrogen reacts with hydrogen to form NH3. Nitrogen is very less reactive and the yield of Ammonia is very less, to increase the percentage yield of Ammonia formed NO is used as a promoter.

What is the significance of autocatalysis?

Auto catalysis is self-catalysis in this process one of the product formed acts as a catalyst and increases the reaction rate.