CATALYSIS

Learning Outcomes
Contents for this page	Related topics
The energy of activation What are catalysts? Types of catalysis Mechanisms of catalysis Additional questions	Reaction rates Chemical equilibrium Le Chatelier's principle	Data Glossary
After studying this section, you will (a) understand what is meant by a catalyst, (b) know the basic principles of homogeneous and heterogeneous catalysis, and (c), have an elementary understanding of the mechanisms involved.

The energy of activation:

This is a revision of the Grade 11 material.

When a reaction proceeds from reactants to products, the change takes place by a complicated route. In the diagram on the right, we will plot the energy change as the reaction proceeds. Before you read on, click here to see how the energy changes as the reaction proceeds.

The overall change in energy is ΔH = P - R, but as the reaction progresses, the energy, H, follows the curve in red. It first rises well above the energy value of the reactants, reaches a maximum, then drops to the energy value for the products.

The "hump" in the curve represents an "energy barrier" of value Ea which is called the ENERGY OF ACTIVATION for the reaction. It may be defined as the minimum energy required to start the chemical reaction. Reactions with high values of Ea take place more slowly than those with small Ea values.

What are catalysts?

For many reversible chemical reactions, the rates at which equilibrium is attained are extremely slow. This is true for most of the important industrial processes, such as the Haber-Bosh process for the synthesis of ammonia from its elements, and the contact process for the preparation of sulphur trioxide.

Substances known as CATALYSTS greatly increase the rates of BOTH the forward AND the reverse reactions, thus shortening the time needed to reach equilibrium. This is achieved by lowering the energy of activation, E_a for the reaction (see the diagram on the left).

It is important to note that at a given temperature, there is no change in the initial energy R of the reactants, nor of the energy P of the products. Thus the value for the energy change in the reaction, ΔH does not change.

In reversible reactions, catalysts DO NOT change the value of the equilibrium constant K_c. The reason for this is that a catalyst will accelerate the rates of the forward and reverse reactions to the same extent. The advantage of using catalysts in reverse reactions resides in the fact that equilibrium is reached much faster than if such catalysts were not present.

Refer now to the Maxwell-Boltzmann distribution curve on the right. Since the energy of activation E_a for the catalyzed reaction is lower than for the uncatalyzed reaction, the proportion of those molecules with sufficient energy to overcome the energy barrier will be greater (compare the areas under the curve in the presence of a catalyst.

At a given temperature therefore, there will be more fruitful collisions and so the rate of the reaction will be increased. In biological systems, almost all reactions are catalysed by certain proteins called ENZYMES. This enables reactions that would normally take place so slowly that changes would be undetectable to occur at rates that are 100 million times or more faster than in the absence of these enzymes. Enzyme catalysis makes life possible!

Catalyst do take part in the reactions which they catalyse, but are regenerated in the process. Thus, a typical catalyst remains unchanged in the overall process. However, certain substances act as "catalyst poisons", and gradually lower the efficacy of catalysts. Platinum metal catalysts, for example, are poisoned by traces of certain sulphur compounds, and have to be replaced at various intervals.

Do you know what "catalytic converters" are all about?

(Click here for a discussion)

Types of catalysis:

There are two types of catalysis. In the case where the reactants and the catalyst are in the same phase, for example, both the reactants and the catalyst are in solution, we have HOMOGENEOUS CATALYSIS. An example of this is the formation of esters from carboxylic acids and alcohols, which is catalysed by acids (indicated by the H⁺ above the reversible arrow in the reaction below):

In the case of HETEROGENEOUS CATALYSIS, the catalyst is usually in the solid phase, while the reactants and products are either all in the gas phase, or all in the liquid phase. A great many important industrial processes, such as the Haber-Bosch process for the industrial production of ammonia, take place by heterogeneous catalysis, where the reactants are gases, and the catalyst (iron, in this case) a solid.

Mechanisms of catalysis:

Chemists are not only interested in reactions as far as reactants and products are concerned, they want to know as precisely as possible what happens to all the atoms during the reaction. It turns out that even simple reactions take place by complex MECHANISMS, usually involving several steps. We will look at two examples, each involving different mechanisms. Generally speaking, catalysts intervene by changing the reaction mechanism, providing pathways that require lower energies of activation, thus increasing the reaction rates.

It is a common misconception, dating back from the days when very little was known about the mechanisms of catalysis, that "catalysts speed up the rates of reactions without taking part in these reactions." You will see below that this is quite wrong.

Homogeneous catalysis:

As an example, look at the hydrolysis of esters by water, which is normally quite slow, and which leads to an equilibrium mixture of acid, alcohol, water and the ester:

In the absence of a catalyst, the bond marked X must be broken by reacting it with water (Step 1), forming the cation C and the anion A. This is a very slow process, with the equilibrium position being very largely on the left. In Step 2, the cation C transfers a hydrogen ion, H⁺ to the anion A, thus forming the products. This step is very fast, with the equilibrium being predominantly on the right of the equation.

The presence of a strong acid, such as sulphuric acid, that provides hydrogen ions, H⁺, alters the mechanism:

A hydrogen ion reacts with the ester to form a cation:
The cation reacts with water to form another cation and the alcohol:
That cation loses H⁺ to form the carboxylic acid. The hydrogen ion is recycled in this step, and is thus not consumed in the overall reaction.

The above reaction mechanisms have been simplified somewhat.

Heterogeneous catalysis:

It is generally accepted that heterogeneous catalysis involving gaseous reactants and solid catalysts involve a number of steps:

	Diffusion of the reactant gases to the catalyst.
	Binding of the reactants to the surface of the catalyst, a phenomenon known as ADSORPTION.
	Breaking of bonds of the adsorbed molecules.
	Making new bonds, forming adsorbed product molecules.
	Detachment of product molecules from the catalyst surface (a process known as DESORPTION) and diffusion away from the catalyst surface.

In the above scheme, catalysis will be favoured if the reactant molecules bind more strongly to the catalyst surface than the product molecules. Ideally, the reactant molecules are concentrated on the catalyst surface, where they undergo far more frequent collisions than if they were in the gas phase. It is suggested that when a reactant is bound to the catalyst surface, the covalent bonds holding its atoms together are distorted and weakened, thus making them more susceptible to breaking.