Page:EB1911 - Volume 06.djvu/44

 order of their freezing-point depressions in equivalent solutions, since the depression of the freezing-point increases with the degree of electrolytic dissociation. Experience confirms this conclusion completely. The degree of dissociation of an acid, at a given concentration, for which its molecular conductivity is, is shown by the theory of electrolytic dissociation to be = /∞; ∞, the molecular conductivity at very great dilution in accordance with the law of Kohlrausch, is u + v, where u and v are the ionic-mobilities (see ). Since u, the ionic-mobility of the hydrogen ion, is generally more than ten times as great as v, the ionic-mobility of the negative acid-radical, ∞ has approximately the same value (generally within less than 10%) for the different acids, and the molecular-conductivity of the acids in equivalent concentration is at least approximately proportional to the degree of electrolytic dissociation, i.e. to the strength.

In general, therefore, the order of conductivities is identical with that in which the acids exert their specific powers. This remarkable parallelism, first perceived by Arrhenius and Ostwald in 1885, was the happy development which led to the discovery of electrolytic dissociation (see ; and ).

Catalysis.—We have already mentioned the fact, early known to chemists, that many reactions proceed with a marked increase of velocity in presence of many foreign substances. With Berzelius we call this phenomenon “catalysis,” by which we understand that general acceleration of reactions which also progress when left to themselves, in the presence of certain bodies which do not change in amount (or only slightly) during the course of the reaction. Acids and bases appear to act catalytically upon all reactions involving consumption or liberation of water, and indeed that action is proportional to the concentration of the hydrogen or hydroxyl-ions. Further, the decomposition of hydrogen peroxide is “catalysed” by iodine-ions, the condensation of two molecules of benzaldehyde to benzoin by cyanogen-ions. One of the earliest known and technically most important instances of catalysis is that of the oxidation of sulphur dioxide to sulphuric acid by oxygen in the presence of oxides of nitrogen. Other well-known and remarkable examples are the catalysis of the combustion of hydrogen and of sulphur dioxide in oxygen by finely-divided platinum. We may also mention the interesting work of Dixon and Baker, which led to the discovery that a large number of gas-reactions, e.g. the combustion of carbon monoxide, the dissociation of sal-ammoniac vapour, and the action of sulphuretted hydrogen upon the salts of heavy metals, cease when water-vapour is absent, or at least proceed with greatly diminished velocity.

“Negative catalysis,” i.e. the retardation of a reaction by addition of some substance, which is occasionally observed, appears to depend upon the destruction of a “positive catalyte” by the body added.

A catalyte can have no influence, however, upon the affinity of a process, since that would be contrary to the second law of thermodynamics, according to which affinity of an isothermal process, which is measured by the maximum work, only depends upon the initial and final states. The effect of a catalyte is therefore limited to the resistances opposing the progress of a reaction, and does not influence its driving-force or affinity. Since the catalyte takes no part in the reaction its presence has no effect on the equilibrium-constant. This, in accordance with the law of mass-action, is the ratio of the separate reaction-velocities in the two contrary directions. A catalyte must therefore always accelerate the reverse-reaction. If the velocity of formation of a body be increased by addition of some substance then its velocity of decomposition must likewise increase. We have an example of this in the well-known fact that the formation, and no less the saponification, of esters, proceeds with increased velocity in the presence of acids, while the observation that in absence of water-vapour neither gaseous ammonium chloride dissociates nor dry ammonia combines with hydrogen chloride becomes clear on the same grounds.

A general theory of catalytic phenomena does not at present exist. The formation of intermediate products by the action of the reacting substance upon the catalyte has often been thought to be the cause of these. These intervening products, whose existence in many cases has been proved, then split up into the catalyte and the reaction-product. Thus chemists have sought to ascribe the influence of oxides of nitrogen on the formation of sulphuric acid to the initial formation of nitrosyl-sulphuric acid, SO2(OH)(NO2), from the mixture of sulphur dioxide, oxides of nitrogen and air, which then reacted with water to form sulphuric and nitrous acids. When the velocity of such intermediate reactions is greater than that of the total change, such an explanation may suffice, but a more certain proof of this theory of catalysis has only been reached in a few cases, though in many others it appears very plausible. Hence it is hardly possible to interpret all catalytic processes on these lines.

In regard to catalysis in heterogeneous systems, especially the hastening of gas-reactions by platinum, it is very probable that it is closely connected with the solution or absorption of the gases on the part of the metal. From the experiments of G. Bredig it seems that colloidal solutions of a metal act like the metal itself. The action of a colloidal-platinum solution on the decomposition of hydrogen peroxide is still sensible even at a dilution of 1/70,000,000 grm.-mol. per litre; indeed the activity of this colloidal-platinum solution calls to mind in many ways that of organic ferments, hence Bredig has called it an “inorganic ferment.” This analogy is especially striking in the change of their activity with time and temperature, and in the possibility, by means of bodies like sulphuretted hydrogen, hydrocyanic acid, &c., which act as strong poisons upon the latter, of “poisoning” the former also, i.e. of rendering it inactive. In the case of the catalytic action of water-vapour upon many processes of combustion already mentioned, a part of the effect is probably due to the circumstance, disclosed by numerous experiments, that the union of hydrogen and oxygen proceeds, between certain temperature limits at least, after the equation H2 + O2 = H2O2, that is, with the preliminary formation of hydrogen peroxide, which then breaks down into water and oxygen, and further, above all, to the fact that this substance results from oxygen and water at high temperatures with great velocity, though indeed only in small quantities.

The view now suggests itself, that, for example, in the combustion of carbon monoxide at moderately high temperatures, the reaction

advances with imperceptible speed, but that on the contrary the two stages

which together result in (I.), proceed rapidly even at moderate temperatures.

Temperature and Reaction-Velocity.—There are few natural constants which undergo so marked a change with temperature as those of the velocities of chemical changes. As a rule a rise of temperature of 10° causes a twofold or threefold rise of reaction-velocity.

If the reaction-coefficient k, in the sense of the equation derived above, viz. k = t&#8202;–1 log {a/(a−x)}, be determined for the inversion of cane-sugar by an acid of given concentration, the following values are obtained:—

here a rise of temperature of only 30° suffices to raise the speed of inversion fifty times.

We possess no adequate explanation of this remarkable temperature influence; but some account of it is given by the molecular theory, according to which the energy of that motion of substances in homogeneous gaseous or liquid systems which constitutes heat increases with the temperature, and hence also the frequency of collision of the reacting substances. When we reflect that the velocity of motion of the molecules of gases, and in all probability those of liquids also, are proportional to the square root of the absolute temperature, and therefore rise by