Chemical Reactions - Rate of Reaction

Catalysts

A catalyst is a substance that speeds up the rate of a chemical reaction without being consumed in the reaction itself. Catalysts do this by providing an alternative reaction pathway with a lower activation energy.

Activation Energy ($E_a$)

The activation energy ($E_a$) is the minimum amount of energy that reactant particles must possess in order for a chemical reaction to occur. It represents the energy barrier that must be overcome for the reaction to proceed.

A lower activation energy means that a larger proportion of reactant particles will have sufficient energy to react, thus increasing the reaction rate.

How Catalysts Work

Catalysts work by forming an intermediate stage in the reaction. This intermediate stage has a lower energy than the uncatalysed reaction pathway. By lowering the activation energy, a greater fraction of molecules can react, leading to a faster reaction rate.

Catalysis and Activation Energy

The key point to remember is that a catalyst decreases the activation energy ($E_a$) of a reaction.

This is represented by the following equation:

$$k = k_0 e^{-E_a/RT}$$

Where:

• k is the rate constant
• k₀ is the initial rate constant
• $E_a$ is the activation energy
• R is the gas constant
• T is the absolute temperature

A lower $E_a$ value will result in a larger value for 'k', indicating a faster reaction rate.

Table Summary

Concept	Description
Catalyst	Substance that speeds up a reaction without being used up.
Activation Energy ($E_a$)	Minimum energy required for a reaction to occur.
Effect of Catalyst	Decreases the activation energy of a reaction.
Rate Constant (k)	A measure of the reaction rate. Lower $E_a$ leads to a higher k.