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Chemical reactions are frequently not fast or selective enough for our practical purposes. Most reactions can lead to various products, which may be useless, toxic, or necessitate tedious cleaning and separation processes. The presence of side products reduces the yield of the synthetic process; slow reactions reduce profitability. Most chemical reactions are a network of parallel and series reactions; some are reversible. The time dependence of chemical processes can be described by chemical kinetics. The energetic aspects, which determine the ratio of products and the temperature dependence of the processes, are determined at the molecular level by quantum chemical rules and at the macroscopic level by thermodynamic functions describing the energy (enthalpy) and the randomness (entropy) of the constituents. In equilibrium, the product ratio is given by the relative free enthalpy of the products (calculated from enthalpy and entropy contributions). However, in several cases, the equilibrium cannot be achieved within practical time limits or is not even recommended. Most reactions from the reactants to the products do not occur spontaneously; the transition state between them has a higher energy level than the reactants or the reaction products. This energy difference is called “activation energy.” It determines the rate of the reaction and its temperature dependence. Catalysts interact with the reactants and reduce the activation energy of the reaction. They release the products and can enter into a new reaction. This cycle results in a considerable acceleration of the reaction, although the concentration of the catalyst is much lower than that of the reactant. As the selectivity, energy consumption, and yield of the chemical industrial processes depend critically on finding suitable catalysts, it is imperative to understand their structure and function in various applications. This module offers a brief introduction to this topic.
Module leader: Bánhegyi György
