The Two Most Common Types Of Catalysts

Catalysts are compounds that accelerate the progress of a reaction without being consumed. They have the ability to increase the rate of reaction in a chemical reaction. Potentially, molecules that would once have taken years to interact, can take seconds with the addition of a catalyst. The overall purpose of a catalyst is to ensure that reactions proceed effectively which is why a range of catalysts are commonly used in many elements of society. Common examples of where catalysts are used include; plastics, clean energy, converting energy sources to fuels, digestion and pharmaceuticals (Hamers 2017). An important feature of a catalyst is that whilst it does increase the rate at which the reaction occurs, the catalyst remains chemically unchanged
The most two common types of catalysts are heterogeneous and homogeneous. The difference between the two is how the catalysts sits in the solution, to be specific, what phase the catalyst is in. A phase in a solution is the visible boundaries between two components. For example, a solution with a liquid and solid has two phases whilst a liquid solution consisting of multiple chemicals has just one phase. Although similar to the physical state, the number of phases in a solution is far more general. For example, a solution of oil and water has two phases because of the visible, separate layers (Clark
Recall that for a reversible reaction, the equilibrium state is one in which the forward and reverse reaction rates are equal. In the presence of a catalyst, both the forward and reverse reaction rates will speed up equally, thereby allowing the system to reach equilibrium faster. However, it is very important to keep in mind that the addition of a catalyst has no effect whatsoever on the final equilibrium position of the reaction. It simply gets it there faster.
Recall that catalysts are compounds that accelerate the progress of a reaction without being consumed. Common examples of catalysts include acid catalysts and enzymes. Catalysts allow reactions to proceed faster through a lower-energy transition state. By lowering the energy of the transition state, which is the rate-limiting step, catalysts reduce the required energy of activation to allow a reaction to proceed and, in the case of a reversible reaction, reach equilibrium more rapidly.
In the presence of a catalyst, the same amounts of reactants and products will be present at equilibrium as there would be in the uncatalyzed reaction. To state this in chemical terms, catalysts affect the kinetics, but not the thermodynamics, of a reaction.