Ionic Dissociation

One of several masters of early physical chemistry was the Swedish chemist Svante August Arrhenius (1859-1927). As a student, Arrhenius turned his attention to electrolytes; that is, to those solutions capable of carrying an electric current.

Michael Faraday (1791-1867) had worked out the laws of electrolysis, and from those laws it had seemed that electricity, like matter, might well exist in the form of tiny particles. Faraday had spoken of ions, which might be considered as particles carrying electricity through a solution. For the next half century neither he nor anyone else had ventured to work seriously on what the nature of those ions might be. This did not mean to no valuable work was done. In 1835, the German physicist Johann Wilhelm Hittorf (1824-1914) pointed out that some ions traveled more rapidly than others. This observation led to the concept of transport number, the rate at which particular ions carried the electric current. But even calculation of this rate still left the nature of ions an open question.

Arrhenius found his entry into the field through the work of the French chemist François Marie Raoult (1830-1901). Like Van't Hoff, Raoult studied solutions. His studies were climaxed in 1887 with his establishment of what is now call Raoult's law: The partial pressure of solvent vapor in equilibrium with a solution is directly proportional to the mole fraction of the solvent.

Without going into the definition of mole fraction, it is sufficient to say that this rule made it possible to estimate the relative number of particles (whether of atoms, molecules or the mysterious ions) of the substance which is dissolved (the solute) and of the liquid in which it is dissolved (the solvent).

In the course of this research, Raoult had measured the freezing points of solutions. Such freezing points were always "depressed"; that is, were lower than the freezing point of the pure solvent. Raoult was able to show that the freezing point was depressed in proportion to the number of particles of solute present in the solution.

But here a problem was created. It was reasonable to suppose that when a substance dissolved in, say, water it broke up into separate molecules. Sure enough, in the case of non-electrolytes such as sugar, the depression of the freezing point fit that assumption. However, when an electrolyte like common salt (NaCl) was dissolved, the depression of the freezing point was twice as great as it should have been. The number of particles present was twice the number of salt molecules. If barium chloride (BaCl₂) was dissolved, the number of particles present was three times as great as the number of molecules.

A molecule of sodium chloride is made up of two atoms, and a molecule of barium chloride is made up of three atoms. It seemed to Arrhenius that when certain molecules were dissolved in a solvent such as water, those molecules broke down into the individual atoms. Furthermore, since those molecules, once broken down, conducted an electric current (whereas molecules such as sugar, which did not break apart, did not carry an electric current), Arrhenius further suggested that the molecules did not break down (or "dissociate") into ordinary atoms, but into atoms carrying an electric charge.

Faraday's ions, Arrhenius proposed, were simply atoms (or groups of atoms) carrying either a positive or a negative electric charge. The ions were either the "atoms of electricity" or they carried those "atoms of electricity". (The latter alternative eventually proved correct.) Arrhenius used his theory of ionic dissociation to account for many facts of electrochemistry.

Arrhenius's ideas, advanced as his Ph.D. thesis in 1884, met with considerable resistance; his thesis was almost rejected. However, Ostwald, impressed, offered Arrhenius a position and encouraged him to continue work in physical chemistry.

In 1889, Arrhenius made another fruitful suggestion. He pointed out that molecules, on colliding, need not react unless they collide with a certain minimum energy, and energy of activation. When this energy of activation is low, reactions proceed quickly and smoothly. A high energy of activation, however, might keep a reaction from proceeding at more than an infinitesimal rate.

If, in the latter case, the temperature were raised so that a number of molecules received the necessary energy of activation, the reaction would then proceed suddenly and quickly, some times with explosive violence. The explosion of a hydrogen-oxygen mixture when the ignition temperature is reached is an example.

Ostwald used this concept profitably in working out his theory of catalysis. He pointed out that the formation of a catalyst-combined intermediate required a smaller energy of activation than the direct formation of the final products required.