What is a catalyst?
A catalyst is defined as a substance which when added to the reacting system increases the rate of a reaction without itself being consumed in the reaction. The phenomenon of increasing rate of reaction with the help of catalyst is known as catalysis and reaction is called as catalyzed reactions.
The catalyst is very important in a chemical industry where reactions are carried out with the help of catalysts.
Classification of catalysts
A catalyst is classified into two kinds depending on whether the catalyst and the reaction mixture are in the same phase or different phases.
A) Homogeneous catalyst
As we say above classification is depends on the phase of the catalyst and reaction mixture. Here the phase of catalyst and reaction mixture is same called as a homogeneous catalyst. Homogeneous catalyst dissolves in the gas phase or solution and works uniformly throughout the gas or solution. Such a phenomenon is called homogeneous catalysis.
We will take an example so that you will understand the above concept.
When aqueous methyl ethyl ester is hydrolyzed with water to give acetic acid and ethyl alcohol, the catalyst used is acid such as HCl. Here both reactants are in the liquid phase as well as a catalyst is in same phase i.e liquid phase.
B) Heterogeneous catalysts
A catalyst which exists in a different phase from the reactants is known as a heterogeneous catalyst. Such a phenomenon is called heterogeneous catalysis.
The heterogeneous catalyst is generally solid and the reactants are either liquids or gases. When the solid catalyst is added to the reacting system, it does not dissolve, unlike homogeneous catalyst. Actually, a reaction occurs on the surface of a solid catalyst. So such kind reactions are also known as surface reactions
The well knows examples of this type of catalysis is hydrogenation of unsaturated compounds. Here catalysts used are platinum, Nickel, palladium etc
Characteristics of Catalysts
According to the definition of catalyst, as it does not consume in reaction, it does not participate. But in reality, it takes part in the reaction. It reacts with one or more molecules to form an intermediate which is a complex of reactant and catalyst.
Reactant + Catalyst → Complex
The complex decomposes to give products and regenerates the catalysts
Complex → Product + Catalyst
So catalyst enters into a chemical reaction, it does not appear in the overall balanced equation.
In a reversible reaction, catalyst increases the rates of both forward and reverse reaction equally. So the position of equilibrium is not influenced and the value of equilibrium constant is not affected by the presence of a catalyst.
An extremely small amount of catalyst causes a considerable increase in the rate of reaction.
Example – One molecule of enzyme catalyze, under certain conditions decomposes 5 million molecules of H2O2 in one minute.
The activation energy of a catalyzed reaction is always lower than that of the same reaction when it is uncatalyzed.
Example – Decomposition of Hydrogen peroxide (H2O2) is slow at room temperature with an activation energy of 76 KJ per mole but when little iodide ion is added, the activation energy reduces to 57 KJ per mole and the rate increases by the factor of 2000.
A catalyst does not affect the energies of reactants and products of the reaction so energies are the same for both catalyzed and uncatalyzed reactions.
In most cases, the catalyst increases the rate of reaction already in progress. It is considered that the catalyst does not initiate the reaction. However, in certain cases, the catalyst is found to initiate the reaction.
Example – A mixture of Hydrogen and Oxygen gas remains unchanged for years at room temperature. If a platinum wire is added to the mixture, immediately the reaction starts and the rate becomes so high that mixture explodes.