Catalysts and Activation Energy

Activation Energy

Activation energy is the minimum amount of energy that reacting particles must have for a chemical reaction to occur. It acts as an energy barrier that reactants need to overcome to be transformed into products.

During a reaction, particles collide. However, not all collisions lead to a reaction. Only those collisions where the particles have energy equal to or greater than the activation energy will result in a successful reaction.

The size of the activation energy affects the reaction rate: a higher activation energy means fewer particles have enough energy to react, so the reaction is slower. Conversely, a lower activation energy means more particles can react, speeding up the reaction.

For example, when lighting a match, the initial spark provides enough energy to overcome the activation energy and start the combustion reaction.

For instance, if a reaction requires an activation energy of 150 kJ/mol, only particles with energy 2150 kJ/mol2 can react. If the temperature increases, more particles reach this energy, increasing the reaction rate.

Catalysts

A catalyst is a substance that speeds up a chemical reaction without being used up or permanently changed in the process. It provides an alternative reaction pathway with a lower activation energy.

Because catalysts lower the activation energy, more particles have enough energy to react when they collide, increasing the rate of reaction.

Catalysts do not affect the overall energy change of the reaction or the position of equilibrium; they only help the reaction reach equilibrium faster.

An everyday example is the catalytic converter in cars, which uses catalysts like platinum and rhodium to speed up the breakdown of harmful gases into less harmful substances.

Effect of Catalysts on Energy Profile

In an energy profile diagram, the activation energy is shown as the peak that reactants must climb to become products.

When a catalyst is used, the peak representing the activation energy is lower, showing that less energy is needed for the reaction to proceed.

The energy levels of the reactants and products remain the same, so the overall energy change of the reaction does not change.

Because the activation energy is lower, the reaction happens faster, as more particles can overcome the energy barrier.

For example, in the decomposition of hydrogen peroxide, manganese dioxide acts as a catalyst, lowering the activation energy and speeding up the reaction without being consumed.

Example: If a reaction requires an activation energy of 100 kJ/mol, only particles with at least this energy can react when they collide.