The addition of a catalyst to a chemical reaction significantly lowers the activation energy, as predicted by quantum chemical calculations. Activation energy is the minimum energy required for a reaction to proceed, and it is often represented as a barrier that reactants must overcome to form products.Catalysts work by providing an alternative reaction pathway with a lower activation energy. They achieve this by forming temporary bonds with the reactants, stabilizing the transition state, and reducing the energy required for the reaction to proceed. This allows the reaction to occur more readily and at a faster rate.Quantum chemical calculations, such as density functional theory DFT and ab initio methods, can be used to model and predict the behavior of molecules and chemical reactions. These calculations can provide insights into the electronic structure, geometry, and energetics of the reactants, transition states, and products involved in a catalyzed reaction.By comparing the activation energies of the uncatalyzed and catalyzed reactions, quantum chemical calculations can help to quantify the effect of a catalyst on the reaction rate. In general, the lower the activation energy, the faster the reaction will proceed, leading to an increased reaction rate. This is consistent with the Arrhenius equation, which relates the reaction rate constant k to the activation energy Ea and temperature T :k = Ae^-Ea/RT where A is the pre-exponential factor, R is the gas constant, and T is the temperature in Kelvin.In summary, the addition of a catalyst to a chemical reaction lowers the activation energy, as predicted by quantum chemical calculations. This leads to an increased reaction rate and improved efficiency of the reaction process.