Catalysis

CATALYSIS, catalysts. Catalyst is a term introduced into science by Berzelius (1835) to denote substances which cause or accelerate chemical processes without taking any visible part in them. Later, Ostwald and his school investigated the question of catalysis from a physico-chemical point of view and laid the foundation for its scientific development. Recently, it has acquired particularly important significance in view of its close connection with the question of enzymes, as well as in view of numerous highly important technical applications. General properties. The main features of the action of catalysts are as follows: 1. Since a typical catalyst does not change during the process, it does not perform work, does not serve as a source of energy, and therefore cannot shift the true chemical equilibrium. In essence, a catalyst only disturbs a false chemical equilibrium (metastable state) or accelerates a process which is accompanied by a decrease in free energy and therefore can occur by itself even without a catalyst, albeit at an immeasurably slow speed. In the opinion of Ostwald, Bredig, and others, a catalyst eliminates "passive resistances" which, for reasons not entirely understood, retard the course of processes; i.e., it acts as a kind of "chemical lubricant." 2. The action of catalysts is for the most part specific, although not to the same degree as in enzymatic processes. Each reaction, and sometimes a specific direction of it, requires its own individual catalyst; thus, for example, the chromic acid ion catalyzes the reaction HBrO₃+HJ, but does not affect the reaction H₂S₂O₈+HJ. The oxidation of naphthalene into phthalic acid occurs only in the presence of a mercury salt, the oxidation of NH₃ by persulfate in the presence of a silver salt, the oxidation of dimethylaniline by oxygen in the presence of a copper salt, and the interaction of oxalic acid and mercuric chloride in the presence of a manganous salt. However, there exist group catalysts that influence series of analogous processes, for example, oxidation reactions. 3. The action of catalysts depends on their available quantity; in particular, soluble catalysts act in dependence on their concentration; in this case, the catalytic effect usually increases to a lesser degree than proportionally to the concentration, i.e., higher concentrations act relatively more weakly than small ones. 4. The absolute quantities of catalysts whose action can still be detected experimentally are for the most part negligibly small. Some soluble catalysts exert a noticeable effect at a content of one billionth of a gram in 1 liter of solution (e.g., copper salts in the reaction of the oxidation of sulfurous acid by oxygen). 5. Frequently, catalysts exhibit an increased conjugated action under the influence of another catalyst (for example, the enhancement of the catalytic effect of iron salts by copper salts in the reaction: H₂S₂O₈+HJ). To such conjugated action, one can attribute the enhancement of the effect of many catalysts under the influence of H⁺ and OH^- ions of the medium, which very often themselves act to accelerate reactions without taking any visible part in them. However, in the majority of cases, the accelerating action of H⁺ and OH^- ions is explained by their formation or disappearance during the process itself, i.e., by direct participation in it (e.g., the reaction: H₂O₂ + 2H⁺+2J^- -> 2H₂O + J₂). Catalysts can be subdivided into two classes: a) soluble catalysts acting on reactions in a homogeneous medium of a solution; b) catalysts acting in a solid insoluble form on reactions occurring in a liquid or, especially often, in a gaseous medium. Since catalysts of type b act upon the contact of substances with their surface, they are often called "contact substances," and catalysis of this kind is called "contact action." A transition between these two groups of catalysts are those which act in a colloidal state; it is to this type that, without a doubt, the vast majority of biochemical agents belong. Homogeneous catalysts (in a homogeneous medium) are extremely numerous and diverse, especially in oxidation reactions of the most varied character and in processes occurring with the participation of water (e.g., the hydrolysis of esters and polysaccharides). Some oxidizers with a high oxidation potential, such as, for example, hydrogen peroxide or persulfates, react extremely slowly, especially in dilute solutions, and the influence of catalysts in minimal doses is manifested here unusually sharply. Such is, for example, the classic reaction: H₂O₂ + HJ, accelerated by the presence of traces of iron salts or molybdic acid (Brodie). In the same way, the hydrolysis (inversion) of cane sugar occurs at a measurable speed only in the presence of H⁺ ions acting catalytically. The theory of such reactions by the majority of authors is reduced to the assumption of the formation during the process of intermediate products which react at a greater speed than the initial substances. To these intermediate substances, one can attribute only greater chemical mobility in comparison with the initial products, but not a higher chemical potential, which is a measure of chemical affinity, since an increase in it would lead to a gain in energy, which would contradict the principle of catalysis. For oxidation reactions, intermediate degrees of oxidation can serve as intermediate substances; but precisely by virtue of the necessary condition that they be chemically mobile, one must look for them not in stable, ordinary degrees of oxidation, but in compounds which, under ordinary conditions, are chemically unstable and therefore cannot be isolated as such. Their formation can for the most part be proven only indirectly; only in exceptional cases is it possible to detect them directly and reduce the total summary reaction to a sequence of several intermediate processes. Such is, for example, the reaction: 2H₂O₂ -> 2H₂O + O₂, accelerated by the presence of KI, for which Bredig and Walton proved the intermediate formation of the compound KIO, which is unstable under ordinary conditions. Frequently, such cases in which intermediate substances are clearly formed are designated by the term "transfer catalysis"; however, there are no sufficient grounds for such a delimitation, since the facts suggest that in other cases as well, catalytic processes have an analogous chemistry at their basis. In many cases, intermediate products are formed by the direct addition of the reacting substances and the catalyst to each other, with the subsequent decomposition of such an aggregate into new molecules and with the regeneration of the catalyst; for example, there are sufficient grounds to assume that during the catalytic influence on oxidation reactions involving hydrogen peroxide, in the first moments, by the addition of hydrogen peroxide to the catalyst, peroxide compounds are formed, which, being unstable, react rapidly with the reducing agent and regenerate the initial form of the catalyst, which again enters into interaction. An analogous chemistry is accepted by Goldschmidt and his school for catalytic reactions in the field of organic substances (esterification, hydrolysis, etc.). Against the theory of intermediate products of catalysis, the consideration was often put forward that such intermediate reactions do not simplify, but complicate the overall process. However, it must be kept in mind that, in general, reactions usually proceed much more complexly than is expressed by an empirical equation; this is indicated by the discrepancy between the order of the reaction calculated from the equation and that observed experimentally (see Chemical Kinetics). Thus, catalysis only replaces some intermediate processes with others that proceed at a greater speed. The state of the catalyst during the process can be imagined as a mobile equilibrium between two chemical states possible for it, whereby it is continuously disturbed and restored by rapid interaction with the components of the reaction mixture, which are thereby introduced into the reaction. The theory of intermediate products finds experimental confirmation in the phenomena of so-called conjugated reactions, or "chemical induction" (Kessler, Luther, Shilov). These phenomena consist in the fact that one of the reactions possible in a given system does not proceed spontaneously, but occurs only in the case when another process is simultaneously proceeding in the same medium. Thus, one can distinguish a primary, spontaneous process and a secondary, induced one. For example, the primary, spontaneous process: HBrO₃+3H₂SO₃ -> HBr+3H₂SO₄; the secondary, induced process: HBrO₃+3H₃AsO₃ -> HBr+3H₃AsO₄. The substance participating only in the primary, spontaneous process bears the name of inductor (H₂SO₃); the substance entering only into the secondary, induced reaction is called the acceptor (H₃AsO₃). The inductor differs from a catalyst in that it itself enters into interaction, changes during the process, and only under this condition draws the acceptor into the reaction, and moreover in a definite ratio, which increases with an increase in the relative concentration of the acceptor.

The maximum amount of acceptor that can enter into a reaction for a certain amount of reacted inductor makes it possible to calculate the composition of the intermediate product linking both conjugated reactions. Thus, for the example given above, one obtains experimentally: HBrO₃ + H₂SO₃ -> HBrO₂ + H₂SO₄ ... 1 O; HBrO₂ + 2H₃AsO₃ -> HBr + 2H₃AsO₄ ... 2 O, i.e., for one atom of oxygen oxidizing the inductor, there are at most two atoms for the acceptor. This points to the intermediate formation of bromous acid, HBrO₂, which, due to its instability, escapes direct observation. Closely related to conjugated reactions are the phenomena of 'autoxidation' or 'activation of O₂': gaseous O₂ oxidizes an acceptor only on the condition of the simultaneous oxidation of an inductor; e.g., O₂ + aldehyde - primary reaction; O₂ + indigo - secondary reaction. For such reactions, the ratio of the inductor and acceptor entering the reaction points to the intermediate formation of peroxide compounds of the inductor by the direct addition of whole molecules of O₂ to it (on this basis, Traube calls such compounds 'holoxides'). These peroxides react secondarily with the acceptor, usually giving it half of the O₂ they have added, which is thus distributed equally between the inductor and the acceptor. In some cases, it is possible to isolate the peroxides in a free state (cerium peroxide, fulvene peroxide) and thereby indisputably prove their formation. Autoxidation reactions were investigated by Schönbein, and later by Traube and Engler; they are of great importance for oxidation reactions involving organic substances (aldehydes, terpenes) and undoubtedly play an important role in biochemical processes. In some cases, a substance acting catalytically is formed during the reaction itself as the final product of the interaction of one of its participants. The action of such a catalyst increases during the reaction as its concentration increases. This phenomenon is called 'autocatalysis'. For example, the manganous ion, formed during the oxidation of oxalic acid by permanganic acid, itself acts catalytically on this same reaction. Such processes are interesting in that, theoretically speaking, they cannot start by themselves without the preliminary addition of an autocatalyst as a seed; but once started, the reaction accelerates itself due to the accumulation of the autocatalyst. The process undergoes a kind of incubation, or a period of growth, similar to many biological processes. Heterogeneous catalysts. Heterogeneous catalysis is of the greatest importance for reactions occurring in a gaseous medium, with metals in a finely divided state (spongy metals) serving primarily as the catalyst: platinum, palladium, nickel, etc., as well as some oxides (aluminum oxide, iron oxide, nickel oxide, zinc oxide) (see Reduction). Solid surfaces of metals act catalytically to a lesser degree; thus, for example, the combination of hydrogen with oxygen occurs in a platinum tube at a low temperature with a measurable speed, while in contact with spongy platinum it quickly leads to an explosion. Heterogeneous catalysis finds wide application in technology. Such are the most important processes: the oxidation of sulfur dioxide into sulfur trioxide in contact with spongy platinum, the synthesis of ammonia from elements in the presence of various catalysts (metallic uranium, iron oxide with certain impurities), the oxidation of ammonia to nitric acid in contact with a platinum mesh. Recently, reactions of hydrogen addition to various organic substances in contact with metallic nickel (Sabatier) have acquired enormous importance; they allow, among other things, the transition from cheap liquid oils to valuable solid fats. Very diverse and interesting reactions are observed in contact with palladium (Zelinsky), as well as with various oxides at high pressure (Ipatiev). By analogous methods, it is possible to technically carry out some organic syntheses, e.g., phosgene from CO and Cl₂, methyl alcohol and liquid fuel (syntol) from CO and H₂. In the preparative practice of organic chemistry, aluminum chloride (Friedel and Crafts) and metallic magnesium, which forms intermediate organomagnesium compounds, play a large role as catalysts; in chlorination reactions, higher inorganic chlorides are used, for example, antimony. Although heterogeneous catalysis can in some cases be experimentally reduced to the formation of more or less unstable intermediate compounds, the phenomenon of adsorption of substances on the surface of solids, especially porous, spongy ones, is of much greater importance for the theoretical explanation of most cases of heterogeneous catalysis. By being adsorbed, the reacting substances come into close contact with each other, and at the same time, due to adsorption forces, their state of polarity is likely changed, which leads to their more energetic interaction. For complex organic molecules, a necessary condition for the acceleration of the process during adsorption is a certain orientation that brings the active groups of atoms closer together (Langmuir, Kruyt). Research further shows that adsorption on the surface of an adsorbent does not occur in a continuous even layer, but in specific areas, and sometimes different ones for different substances, the participation of which can be gradually eliminated thanks to this (topochemical studies of Pease, Taylor, etc.). Similar to the conjugated action of soluble catalysts, heterogeneous catalysts are also often enhanced in their action due to certain specific impurities (the influence of promoters). It is very important to note that heterogeneous catalysts are strongly subject to the phenomena of 'fatigue' and 'poisoning', which must be linked to the irreversible adsorption of foreign substances and final or side products of the reaction itself on the contact surface, which thereby loses, completely or partially, its ability to adsorb reacting substances, i.e., its catalytic activity. In the classic example of the oxidation of sulfur dioxide by oxygen, Bodenstein proved experimentally and mathematically that the speed of the process depends ultimately on the speed of the diffusion of the two reacting gases to the surface of the platinum through a layer of the formed sulfur trioxide, while the act of combination of the adsorbed SO₂ and O₂ itself occurs practically instantaneously. Many substances capable of 'poisoning' heterogeneous catalysts are also typical poisons for the organism (hydrocyanic acid, H₂S, mercury salts, etc.). Such 'poisoning' is also observed for catalysts of a colloidal nature, for which, as for typical heterogeneous catalysts, the phenomenon of adsorption of substances on the surface of colloidal particles is characteristic (Bredig). From this point of view, the catalytic action of colloidal solutions of platinum, palladium, and other metals has been studied in detail. It is necessary to mention that in reactions in a gaseous medium, the catalytic action of water vapor (humidity) is often observed in cases where water does not participate directly in the chemistry of the process; absolutely dry gases enter into interaction with each other with difficulty (Baker, etc.). The walls of vessels, even glass ones, also exert a catalytic influence on gas reactions, as adsorption is observed on them as well. In conclusion, it can be pointed out that in addition to the accelerating action of substances on reactions, a retarding action is sometimes observed-negative catalysis. For homogeneous catalysis, it reduces in the vast majority of cases to the elimination of the positive catalyst by binding it in the form of an inactive compound (e.g., the influence of organic complexing agents on the catalytic action of copper ions). The weakening of heterogeneous catalysis is caused by the retarding influence on those adsorption phenomena which under normal conditions lead to catalysis (this includes the above-mentioned cases of 'poisoning' of heterogeneous and colloidal catalysts). It is possible, however, to imagine a slowing down of a reaction also by changing its chemistry under the influence of a catalyst.