A diminution in the number of positive ions would thus naturally be accompanied by a rise in potential gradient.
It is clear that, when two opposite streams of ions move past each other, equivalent quantities are liberated at the two ends of the system.
In 1887 Svante Arrhenius, professor of physics at Stockholm, put forward a new theory which supposed that the freedom of the opposite ions from each other was not a mere momentary freedom at the instants of molecular collision, but a more or less permanent freedom, the ions moving independently of each other through the liquid.
Thus the osmotic pressure, or the depression of the freezing point of a solution of potassium chloride should, at extreme dilution, be twice the normal value, but of a solution of sulphuric acid three times that value, since the potassium salt contains two ions and the acid three.
The rate of loss of charge is thus largely dependent on the extent to which ions are present in the surrounding air.
It depends, however, in addition on the natural mobility of the ions, and also on the opportunities for convection.
The volume of air from which the ions have been extracted being known, a measure is obtained of the total charge on the ions, whether positive or negative.
The conditions must, of course, be such as to secure that no ions shall escape, otherwise there is an underestimate.
I + is used to denote the charge on positive ions, I_ that on negative ions.
Thus, associ- 1~e] ions of Agropyrum (Triticum) junceum, of Carex arenaria, of ~ ~nmophila (Psamma) arenaria, and of other plants occur on sa rid dunes: the associations are related by the general identity ph the habitat conditions, namely, the physiological dryness f d the loose soil; but they are separated by differences in f~1
A study of the products of decomposition does not necessarily lead directly to a knowledge of the ions actually employed in carrying the current through the electrolyte.
Since the electric forces are active throughout the whole solution, all the ions must come under its influence and therefore move, but their separation from the electrodes is determined by the electromotive force needed to liberate them.
When the amount of this ion in the surface layer becomes too small to carry all the current across the junction, other ions must also be used, and either they or their secondary products will appear also at the electrode.
In aqueous solutions, for instance, a few hydrogen (H) and hydroxyl (OH) ions derived from the water are always present, and will be liberated if the other ions require a higher decomposition voltage and the current be kept so small that hydrogen and hydroxyl ions can be formed fast enough to carry all the current across the junction between solution and electrode.
When the ions are set free at the electrodes, they may unite with the substance of the electrode or with some constituent of the solution to form secondary products.
At the electrodes, however, the small quantity of hydrogen and hydroxyl ions from the water are liberated first in cases where the ions of the salt have a higher decomposition voltage.
If the current be so strong that new hydrogen and hydroxyl ions cannot be formed in time, other substances are liberated; in a solution of sulphuric acid a strong current will evolve sulphur dioxide, the more readily as the concentration of the solution is increased.
Here the ions are potassium and the group Ag(CN)2.1 Each potassium ion as it reaches the cathode precipitates silver by reacting with the solution in accordance with the chemical equation K--+KAg(CN) 2 =2KCN+Ag, while the anion Ag(CN) 2 dissolves an atom of silver from the anode, and re-forms the complex cyanide KAg(CN) 2 by combining with the 2KCN produced in the reaction described in the equation.
The salt must therefore be derived from an acid, chloroplatinic acid, H 2 PtC1 6, and have the formula Na 2 PtC1 6, the ions being Na and PtCls", for if it were a double salt it would decompose as a mixture of sodium chloride and platinum chloride and both metals would go to the cathode.
The opposite parts of an electrolyte, which work their way through the liquid under the action of the electric forces, were named by Faraday the ions - the travellers.
Hittorf (1853) to the unequal speeds with which he supposed the two opposite ions to travel.
If the ions move at equal rates, the salt which is decomposed to supply the ions liberated must be taken equally from the neighbourhood of the two electrodes.
If the black ions move twice as fast as the white ones, the state of things after the passage of a current will be represented by the lower part of the figure.
Here the middle part of the solution is unaltered and the number of ions liberated is the same at either end, but the amount of salt left at one end is less than that at the other.
There is reason to believe that in certain cases such complex ions do exist, and interfere with the results of the differing ionic velocities.
If some of the anions, instead of being simple iodine ions represented chemically by the symbol I, are complex structures formed by the union of iodine with unaltered cadmium iodide - structures represented by some such chemical formula as I(CdI 2), the concentration of the solution round the anode would be increased by the passage of an electric current, and the phenomena observed would be explained.
It is found that, in such cases as this, where it seems necessary to imagine the existence of complex ions, the transport number changes rapidly as the concentration of the original solution is changed.
For instance, to take the two solutions to which we have already referred, we have of ions between molecules at the instants of molecular collision only; during the rest of the life of the ions they were regarded as linked to each other to form electrically neutral molecules.
In such salts as potassium chloride the ions seem to be simple throughout" a wide range of concentration since the transport numbers for the same series of concentrations as those used above run Potassium chloride 0.5 1 5, 0.515, 0.514, 0.513, 0.509, 0.508, 0.507, 0.507, 0.506.
Kohlrausch formulated a theory of electrolytic conduction based on the idea that, under the action of the electric forces, the oppositely charged ions moved in opposite directions through the liquid, carrying their charges with them.
On the view of the process of conduction described above, the amount of electricity conveyed per second is measured by the product of the number of ions, known from the concentration of the solution, the charge carried by each of them, and the velocity with which, on the average, they move through the liquid.
If a solution, let us say of sugar, be confined in a closed vessel through the walls of It is probable that in both these solutions complex ions exist at fairly high concentrations, but gradually gets less in number and finally disappear as the dilution is increased.
The concentration is known, and the conductivity can be measured experimentally; thus the average velocity with which the ions move past each other under the existent electromotive force can be estimated.
The velocity with which the ions move past each other is equal to the sum of their individual velocities, which can therefore be calculated.
Hence the absolute velocities of the two ions can be determined, and we can calculate the actual speed with which a certain ion moves through a given liquid under the action of a given potential gradient or electromotive force.
The details of the calculation are given in the article Electric conduction, § where also will be found an account of the methods which have been used to measure the velocities of many ions by direct visual observation.
The results go to show that, where the existence of complex ions is not indicated by varying transport numbers, the observed velocities agree with those calculated on Kohlrausch's theory.
The verification of Kohlrausch's theory of ionic velocity verifies also the view of electrolysis which regards the electric current as due to streams of ions moving in opposite directions through the liquid and carrying their opposite electric charges with them.
There remains the question how the necessary migratory freedom of the ions is secured.
As we have seen, Grotthus imagined that it was the electric forces which sheared the ions past each other and loosened the chemical bonds holding the opposite parts of each dissolved molecule together.
Arrhenius pointed out that these exceptions would be brought into line if the ions of electrolytes were imagined to be separate entities each capable of producing its own pressure effects just as would an ordinary dissolved molecule.
(I) In very dilute solutions of simple substances, where only one kind of dissociation is possible and the dissociation of the ions is complete, the number of pressure-producing particles necessary to produce the observed osmotic effects should be equal to the number of ions given by a molecule of the salt as shown by its electrical properties.