CATALYSIS. - Certain substances increase the speed of certain chemical reactions when they are present in the environment in which these reactions occur: these are called catalysts, and their action is termed catalysis. It was Berzelius who proposed this term, deriving it from the Greek: καταλυσειν = to unbind, to dissolve.
The following criteria allow one to distinguish catalysis from changes in speed due to other causes: a) at the end of the reaction the catalyst is found chemically unchanged; the word “chemically” is used because sometimes, in the case of heterogeneous catalysis, having taken part in the reaction, it is found in a physical state notably different from the initial one; b) a small quantity of catalyst is often sufficient to transform large quantities of substances; c) the catalyst does not alter the final state of equilibrium of a reversible reaction.
A decrease in speed is called negative catalysis. Sometimes it seems that the catalyst provokes reactions that would not take place without it: even in such cases, however, it is admitted that the catalyst does not provoke but only accelerates a reaction already in progress, albeit extremely slow, or, more precisely, facilitates the transformation of a system in false equilibrium.
Homogeneous catalysis occurs when the substances that react and the catalyst belong to the same phase: that is, they are all gaseous or dissolved in water or another solvent; heterogeneous catalysis occurs when the solid catalyst is immersed in a mixture of gases or in a liquid where the reaction takes place. Examples of homogeneous catalysis are those provided by hydrogen ions (v. ACQUA), which catalyze the inversion of sucrose and the hydrolysis of esters into their organic acid and alcohol components. More important is heterogeneous catalysis, examples of which include: the formation of water from hydrogen and oxygen gases under the action of “platinum sponge,” the formation of ammonia from hydrogen and nitrogen gases in the presence of iron powder, and the combination of sulfur dioxide with oxygen to give sulfur trioxide in the presence of platinum asbestos.
The mechanism by which the catalyst increases the speed of reaction is very often very complex and poorly understood: the simple schemes with which the chemist habitually represents reactions highlight only the reacting products, those that are formed, and their relative proportions; but, with or without a catalyst, reactions are much less simple. In recent years (Hinshelwood, 1946), for example, to explain all the particulars of the reaction 2H₂ + O₂ = 2H₂O, it has been necessary to admit numerous intermediate reactions in which molecules such as H, OH, O, HO₂ appear—molecules unknown to elementary chemistry. The speed of these intermediate reactions is also influenced by molecules of other gases apparently inert that may be present, as well as by the walls of the vessel in which the behavior of the gaseous mixture is studied.
Biological chemistry (v.) offers numerous examples of catalysis. Some of these are specific: the catalyst, which in this case is called an enzyme, favors only a single, very particular reaction or a restricted group of reactions; others, however, are nonspecific, in that the same catalyst favors various reactions, even of totally different types. Substances in the colloidal state, both inorganic and organic, generally have catalytic properties; and one of the reasons why the chemist is unable to reproduce many reactions of living organisms lies precisely in the difficulty of creating an artificial catalytic colloidal environment similar to that of cells.