What is the Difference Between Catalyst and Intermediate

The key difference between catalyst and intermediate is that a catalyst is useful at the beginning of the reaction and is regenerated at the end, whereas an intermediate is formed during the chemical reaction and does not exist at the end of the reaction.

The terms catalyst and intermediate are very important in chemical reactions. A catalyst is a chemical compound that can increase the rate of a reaction without itself being consumed, whereas an intermediate is a molecule that forms from two or more reactants and undergoes further reaction to give final products.

CONTENTS

1. Overview and Key Difference
2. What is a Catalyst  
3. What is an Intermediate 
4. Catalyst vs Intermediate in Tabular Form
5. Summary – Catalyst vs Intermediate

What is a Catalyst?

A catalyst is a chemical compound that can increase the rate of a reaction without itself being consumed. Therefore, this compound can continue to act repeatedly. Due to this reason, only a small amount of catalyst is required for a certain chemical reaction.

A catalyst provides an alternative pathway for a chemical reaction by reducing the activation energy of a reaction. Here, the catalyst combines with the reactant to create an intermediate product, and after the completion of the required reaction, the catalyst leaves the intermediate and regenerates.

There are two types of catalysts; they are homogeneous and heterogeneous catalysts. In homogeneous catalysts, the molecules are in the same phase as reactant molecules. However, in heterogeneous catalysts, the molecules are in a different phase to that of reactant molecules. Enzymes are a good example of biological catalysts.

An intermediate is a molecule that forms from two or more reactants and undergoes further reactions to give final products. An intermediate is formed in multiple-step reactions. Most of the time, complicated chemical reactions require more than one step to complete the reaction to get the desired final product. In these reactions, all the reaction steps other than the last step give intermediates; the last step gives the product rather than giving an intermediate. Therefore, an intermediate is unstable, and it tends to quickly undergo further reaction.

Typically, intermediates occur in the reaction mixture very rarely due to their unstable nature. They exist for a short time. Moreover, it is very difficult to isolate an intermediate because it tends to react further. In a particular reaction, there can be a very high number of intermediate molecules in every reaction step. Sometimes, it is very difficult to identify these molecules.

We can distinguish between intermediates and molecular vibrations. These typically have similar lifetimes and are merely transitions. Usually, these molecules are highly reactive. A good example would be the esterification of a diol, where a monoester is produced in the first step, and a dioester is formed in the second (final) step.

The key difference between catalyst and intermediate is that a catalyst is useful at the beginning of the reaction and is regenerated at the end, whereas an intermediate is formed during the chemical reaction and does not exist at the end of the reaction. Moreover, catalysts are stable, while intermediates are highly unstable.

The below infographic presents the differences between catalyst and intermediate in tabular form for side-by-side comparison.

Summary – Catalyst vs Intermediate

The key difference between catalyst and intermediate is that a catalyst is added at the beginning of the reaction and regenerated at the end of the reaction whereas an intermediate is formed during the reaction and is not regenerated at the end of the reaction.

Reference:

1. Helmenstine, Anne Marie, Ph.D. “What Is a Reaction Intermediate?” ThoughtCo, Aug. 27, 2020.

Image Courtesy:

1. “Catalyst Energy Diagram” By Emma Ambrogi – Own work (CC BY-SA 4.0) via Commons Wikimedia
2. “Hess Cycle Diagram” By SGDWN – Own work (CC BY-SA 4.0) via Commons Wikimedia

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