6.8.5. Effect of a Catalyst

A catalyst increases the rate of the chemical reaction by making available a new low energy pathway for the conversion of reactants to products.

Catalysts are substances that accelerate chemical reactions but are not consumed or changed during the reaction. They function by providing an alternative reaction pathway that has a lower activation energy than the uncatalyzed reaction. This means that more reactant molecules can overcome the energy barrier, leading to an increased rate of reaction.

Imagine you are hiking a steep mountain. If there is a well-defined path that winds around the mountain, you can reach the top more easily than if you take a direct, steep route. The well-defined path is similar to what a catalyst does in a chemical reaction—it provides a more efficient route for the reactants to convert into products.

It increases the rate of forward and reverse reactions that pass through the same transition state and does not affect equilibrium.

By lowering the activation energy, a catalyst speeds up both the forward and reverse reactions equally. This means that while the reaction rate increases, the position of the equilibrium remains unchanged. The catalyst does not alter the equilibrium composition—it only allows the system to reach equilibrium faster.

Think of a racing circuit where both cars are racing in opposite directions. If you set up a shortcut for both cars, they both would reach their destination quicker, but the results of the race (who wins) remain unchanged. In the same way, a catalyst allows the reactions to occur faster without changing the final state of the reaction mixtures.

Catalyst lowers the activation energy for the forward and reverse reactions by exactly the same amount.

Regardless of the direction of the reaction, a catalyst facilitates both paths similarly by providing a stable intermediate (transition state) that lowers the energy barrier. This property ensures that neither direction is favored, preserving the ratio of reactants to products at equilibrium, even if the rate of reaching that state changes.

Consider a bridge that connects two sides of a river. Whether you are going to the left or to the right, using the bridge makes the crossing much easier and faster, without changing how many people are going in either direction. That's what a catalyst does—facilitates the journey without changing the destination.

German chemist, Fritz Haber discovered that a catalyst consisting of iron catalyzes the reaction to occur at a satisfactory rate at temperatures, where the equilibrium concentration of NH3 is reasonably favourable.

The Haber Process for ammonia synthesis is a practical application of catalysts in industry. The reaction between nitrogen and hydrogen to form ammonia is exothermic and slow without a catalyst. By introducing iron as a catalyst, the reaction rate increases significantly, allowing for higher production of ammonia under economically feasible conditions—around 500°C and 200 atm.

Consider a factory assembly line. If you have machines (catalysts) that speed up each step, the entire process of assembling products speeds up, allowing you to produce more items in the same amount of time without changing the design of the items. This analogy reflects how catalysts enhance production efficiency in chemical manufacturing.

Catalyst does not affect the equilibrium composition of a reaction mixture.

While catalysts are effective at speeding up chemical reactions, they do not change the final concentrations of reactants and products at equilibrium. This is an important limitation, as it means that while catalysts can increase the speed of reaching equilibrium, they cannot create more products than would be available without them.

Imagine setting a timer for a certain amount of time when baking a cake. The timer helps you manage your time better, allowing the cake to be done faster, but it won’t change the recipe or how much cake you ultimately have. Similarly, catalysts enhance the speed but do not change the yield of a reaction.

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