CATALYSIS. — Certain substances increase the rate of certain chemical reactions when they are present in the environment in which these reactions take place: they are called catalysts, and their action catalysis. Berzelius was the one to propose this term, deriving it from the Greek: καταλύειν = to loosen, dissolve.
The following criteria make it possible to distinguish catalysis from variations in rate due to other causes: a) at the end of the reaction the catalyst is found to be chemically unaltered: “chemically” is specified because sometimes, in the case of heterogeneous catalysis, having taken part in the reaction, it is found in a physical state notably different from its 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 rate is called negative catalysis. Sometimes it seems that the catalyst causes reactions that would not take place without it: even in such cases, however, it is assumed that the catalyst does not cause but merely accelerates a reaction already under way, even if extremely slowly or, more exactly, facilitates the transformation of a system in false equilibrium.
Homogeneous catalysis occurs when the reacting substances 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 in which the reaction takes place. Examples of homogeneous catalysis are those provided by hydrogen ions (v. arqua), which catalyze the inversion of sucrose and the hydrolysis of saline ethers into the organic acid and alcohol components. More important is heterogeneous catalysis, examples of which are: 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 produce sulfur trioxide in the presence of platinized asbestos.
The mechanism by which the catalyst increases the rate of reaction is for the most part very complex and poorly understood: the simple schemes with which the chemist habitually represents reactions bring out only the products that react, those that are formed, and the relative proportions; but, with or without a catalyst, reactions are far less simple. In recent years (Hinshelwood, 1946), for example, in order to explain all the features of the reaction 2H₂ + O₂ = 2H₂O, it has been necessary to posit numerous intermediate reactions in which unstable molecules such as H, OH, O, and HO₂ appear, unknown to elementary chemistry. The rate of these intermediate reactions is also influenced by molecules of other apparently inert gases that may be present, as well as by the walls of the vessel in which the behavior of the gas mixture is studied.
chemistry (v.) offers numerous examples of catalysis. Some of them are specific: the catalyst, which in this case is called an enzyme, promotes one highly specific reaction or a limited group of reactions; others, on the other hand, are nonspecific, since the same catalyst promotes various reactions, even those of entirely different types. Substances in the colloidal state, both inorganic and organic, generally possess 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.