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Catalysts and Speeding up Reactions

Before a chemical reaction can take place, reacting molecules must crash into one another with enough force to disrupt their original bonds, enter a higher energy transition state, and then form the new bonds of the product molecules. If they just bump into each other, they will just bounce off and nothing will happen. No chemical reaction will take place.

Even in the case of chemical reactions that obey the second law of thermodynamics, such as the reaction between hydrogen gas and oxygen gas to produce water, the reactants simply bump around and off each other and no observable chemical reaction happens. This is very frustrating if you are a human chemist waiting for the product to form in your test tube, or a yeast cell waiting for the sugars in the grape juice to break down. Naturally, both organisms would like to find ways of speeding up chemical reactions so that a reasonable amount of product is formed in a reasonable period of time.

There are four ways of speeding up a chemical reaction.

  1. Extra energy can be added to the reacting molecules so that their kinetic energy is increased, they move faster, crash together with more force and thus produce the required products. Heating the reaction mixture is a common and convenient way of doing this.

  2. More reactant molecules can be brought into closer contact with one another. If there are more and more molecules, closer and closer together there is a better and better chance of a productive collision. Increasing the concentration of the reactants increases the collisions and increases the number of reactions. More product will form.

  3. Those product molecules that do form can be removed. This prevents them from re-forming the original molecules (thus defeating the purpose of the reaction). In strongly exergonic reactions this may not help much, but in mildly exergonic reactions, decreasing the concentration of products will 'pull' a reaction in the direction of more product molecules.

  4. The amount of extra energy, the activation energy needed to get a reaction started, can be reduced. If less activation energy is needed, then more reactant molecules will have enough energy to make productive collisions, and the speed of the reaction will increase. A catalyst is an agent that lowers the activation energy of a reaction. In the presence of a catalyst, therefore, the speed of the reaction is increased.

Many pure metals, like mercury or platinum make good inorganic catalysts. These agents provide large surfaces on which the reactant molecules are absorbed and oriented towards each other. Held in just the right relationship, the reactant molecules need less activation energy to make the collision productive, so the chemical reaction can take place faster at lower temperatures. Catalysts are use a lot in the manufacture of everything from rubber to margarine, and most modern cars have 'catalytic converters' as part of their exhaust systems. Dangerous chemicals produced in the car's engine pass over a platinum catalyst before being allowed into the air. On the surface of the metal the dangerous chemicals react and become less harmful.

Catalysts lower the activation energy of a chemical reaction, increase the rate (speed) of the reaction without themselves becoming permanently altered.

Science@a Distance
© 2001, Professor John Blamire