Selectivity And Activity Explained

A catalyst is a substance that increases the rate of a reaction without changing itself in terms of mass and composition. Take this chemical reaction, for example:

2KClO₃2KCl+3O₂. 

In this reaction, potassium chlorate is decomposed into potassium chloride and oxygen. This is a slow reaction that occurs in the temperature range of 653-873K. If we add manganese dioxide to this reaction, it takes place at an accelerated rate and in the temperature range of 473-633K. Hence manganese dioxide acts as a catalyst in the above reaction.

The activity and selectivity of a catalyst are two of the most important characteristics. A catalyst transforms a chemical reaction. Most industrial syntheses and biological reactions require catalysts. Enzymes are a common catalyst present in nature. 

Selectivity of A Catalyst

A catalyst is selective in nature, which means it directs the reaction such that a particular product becomes a major product or in higher yield over other products. In other words, a catalyst can accelerate a particular reaction selectively and suppress other side reactions.

In the reactions given below, the reactants are the same (CO and H2) but we get different products with different catalysts.

CO+3H2CH4+H2O in the presence of nickel as the catalyst.

CO+H2HCHO in the presence of copper as the catalyst.

A substance may act as a catalyst in a particular reaction but may not catalyse another reaction at all. 

Activity 0f A Catalyst

The activity of a catalyst tells us how fast a reaction will proceed in the presence of a catalyst. This is an important characteristic of a catalyst as a catalyst is chosen based on its activity. It describes the accelerating power of the catalyst on the rate of reaction.

Some substances increase the activity of a catalyst. These are known as promoters. In Haber’s process, molybdenum acts as a promoter by increasing the activity of iron (catalyst).

Some substances decrease the activity of a catalyst. These are known as poisons.

The activity of a solid catalyst depends on how strongly gas molecules or atoms form chemical bonds with the solid surface of the catalyst (also known as chemisorption). These bonds can be ionic or covalent. The reactants have to get adsorbed on the surface of the catalyst to get activated. However, the adsorption strength must not be too strong and other reactants should also get space on the catalyst’s surface to get adsorbed. 

Importance of Catalysts

A total of 85-90% of products in the chemical industry are synthesised using catalysts. Catalysts are indispensable to the chemicals industry. 

All biological processes require enzymes (also a catalyst). Some examples of catalytic processes by enzymes are the fermentation of sugar to ethanol and the conversion of ethanol to acetic acid.

The presence or absence of catalysts highly affects the reaction speed. A catalyst increases the reaction rate by lowering the activation energy needed for a reaction to occur. A catalyst does not change itself or destroy itself by presenting into a reaction. It just acts neutral to speed up the speeds. For example, H₂ and O₂ do not naturally combine. It combines in the presence of a small quality of platinum. This platinum acts as a catalyst to speed up the reaction rate. Without catalysts, many reactions may not be economical and many would not even occur. 

Homogeneous And Heterogeneous Catalysis 

Catalytic reactions can be broadly classified into homogeneous and heterogeneous catalysis based on the phase in which the reaction occurs.

When the reactants and products are present in the same phase (liquid or gas), it is known as homogeneous catalysis.

For example, in the reaction 2SO₂ (g) + O₂ (g) 2SO₃ (g), in presence of NO(g) as the catalyst, the reactants, catalyst and products are in the gas phase.

In the reaction CH₃COOCH₃ (l) + H₂O(l) →CH₃COOH(aq) + CH₃OH(aq) in presence of HCl as catalyst ,the reactants and catalyst are in the liquid phase.

In heterogeneous catalysis, the reactants and catalysts are in different phases.

In the reaction below, sulphur dioxide is oxidised to sulphur trioxide in the presence of platinum (catalyst). 

2SO₂(g) + O₂   2SO₃ (g)

The reactant is in the gas phase whereas platinum is in solid. Therefore this is a heterogeneous catalysis.

In Haber’s process, nitrogen and hydrogen are converted to ammonia in the presence of iron. 

N2(g)+3H2(g)2NH3(g)

Here the reactants are in the gas phase whereas Fe (catalyst) is in the solid state.

Hydrogenation of vegetable oils to ghee with nickel as a catalyst is also an example of heterogeneous catalysis.

Vegetable oils (l) + H2 (g)  Vegetable ghee(s)

The vegetable oils are liquid, hydrogen is gas and Nickel (catalyst) is solid.

Conclusion

Catalysts are substances that increase the rate of reaction but do not get consumed themselves. Their mass and composition remain almost unaltered after the reaction.

Catalysts increase the rate of the reaction by lowering the activation energy. They do not affect the equilibrium position in a reaction.

Catalysis is majorly of two types: homogeneous and heterogeneous.

Activity and selectivity are important characteristics of a catalyst. Activity determines how fast the rate of reaction will occur whereas selectivity is the ability of the catalyst to speed up a particular reaction while suppressing other side reactions.

Enzymes are the catalysts of nature.

Without catalysts, most reactions in the chemical industry would not be possible or economical.