Chemical Change: Catalysts

Catalysts are substances that speed up a chemical reaction, but are not consumed by it. After the reaction, they are available to participate again, so little quantity is required to be effective.
Catalysts usually work by offering an alternative pathway or mechanism for a reaction that requires less activation energy.
Activation energy is the initial input of energy required to start a reaction. By decreasing this, more reactants have enough energy to react, and the rate of reaction is increased.
The place where the reaction occurs is typically on the catalyst’s surface, which provides a platform where reactant molecules can come together. This reduces the energy they need to react.
In industry, catalysts play an essential role, particularly in the production of ammonia in the Haber process where an iron catalyst is used, and in catalytic converters in cars where a platinum-rhodium catalyst is used to decrease harmful emissions.
Enzymes, which are responsible for regulating almost all biochemical reactions in living cells, are natural biological catalysts made of proteins. For example, enzymes in the digestive system break down food.
Catalytic poisoning can occur if a substance binds to the catalyst and blocks the active sites where the reaction takes place, impairing the reaction.
Catalysts can be either homogeneous, existing in the same phase as the reactants, or heterogeneous, existing in a different phase. For example, if a catalyst is a solid and the reactants are gases, the catalyst is heterogeneous.
Remember, catalysts don’t affect the position of the equilibrium in a reversible reaction. They only speed up the rate at which equilibrium is reached.