7.6 Catalysis

Many industrial reactions require high temperatures to achieve a fast reaction rate, which is necessary to maximize product yield in a given time. However, high-temperature processes can pose safety risks and are not suitable for chemical species that are unstable at high temperatures. An alternative method to increase reaction rates is therefore highly valuable.

This alternative is catalysis. By introducing a catalyst, we can change the reaction mechanism to one that has a lower activation energy, thereby increasing the reaction rate.

Catalyst: A substance that accelerates a chemical reaction but remains chemically unchanged at the end of the reaction.

Catalysis: The process of increasing the rate of a chemical reaction by adding a catalyst.

A catalyst provides a new, alternative pathway for the reaction with a lower activation energy ($E_a$). As shown in the energy profile diagram, this lower energy barrier allows more reactant molecules to have sufficient energy to overcome it, leading to a faster reaction rate.

Key Points:

• A catalyst increases the rate of reaction by decreasing its activation energy.
• A catalyst has no effect on the overall thermodynamics or enthalpy ($\Delta H$) of the reaction. It does not change the energy of the reactants or products.
• A catalyst cannot initiate a reaction that is not thermodynamically favorable.

Types of Catalysis

In the FSc curriculum, catalysis is generally classified based on the physical state of the reactants and the catalyst:

1. Homogeneous Catalysis: If the catalyst and the reactants are in the same phase (e.g., all gases or all in solution).
2. Heterogeneous Catalysis: If the catalyst and the reactants are in different phases (e.g., a solid catalyst with gaseous reactants).

Autocatalysis and Catalytic Poisoning

Sometimes, a product formed during the reaction acts as a catalyst for that same reaction. This phenomenon is known as Autocatalysis.

Conversely, certain substances can decrease or destroy the activity of a catalyst. This is known as Catalytic Poisoning.

Example: Catalytic Destruction of Ozone in the Stratosphere

The conversion of ozone (${\mathrm{O}}_3$) to molecular oxygen (${\mathrm{O}}_2$) by an oxygen atom ($\mathrm{O}$) in the stratosphere is a naturally occurring process. However, this direct reaction has a relatively high activation energy.

Uncatalyzed Reaction: The direct reaction between ozone and an oxygen atom is slow due to its high activation energy. $O_3+O\to 2O_2(E_a=17.1{\text{kJ mol}}^{-1})$

Catalyzed Reaction: Chlorofluorocarbons (CFCs) from human activities can diffuse into the stratosphere. There, they absorb high-energy ultraviolet (UV) light, which breaks the carbon-chlorine bonds and releases highly reactive chlorine atoms ($\mathrm{Cl}$). These chlorine atoms act as catalysts for ozone destruction.

The chlorine atom provides a new, two-step mechanism with a much lower overall activation energy:

1. Step 1: A chlorine atom reacts with an ozone molecule. $O_3+Cl\to O_2+ClO(E_a=2.1{\text{kJ mol}}^{-1})$
2. Step 2: The chlorine monoxide ($\mathrm{ClO}$) intermediate reacts with an oxygen atom, regenerating the chlorine atom. $O+ClO\to O_2+Cl(E_a=0.4{\text{kJ mol}}^{-1})$

Net Reaction: By adding the two steps, we see that the chlorine atom is regenerated and does not appear in the overall equation. The net result is the same as the uncatalyzed reaction, but the pathway is different and much faster. $O_3+O\to 2O_2(\text{Overall }{E}_a\approx 2.5{\text{kJ mol}}^{-1})$

The chlorine-catalyzed reaction has a substantially lower activation energy ($2.5{\text{kJ mol}}^{-1}$) than the direct reaction ($17.1{\text{kJ mol}}^{-1}$), which is why a small amount of chlorine can destroy a large amount of ozone.

Enzyme Catalysis

Enzymes are biological catalysts. They are complex protein molecules with high molecular weights and are extremely specific in their action.