Dissociation Electrolytic Dissociation or Ionisation
Dissociation
A certain number of substances are known which apparently do not conform to the laws which have been explained in the last chapter. For example, the compound of ammonia with hydrochloric acid, which has the formula NH 4 C1, should have the density 26.75, f r the atomic weights of the elements it contains are N= 14 ; H= i ; Cl= 35.5 ; and the molecular weight is the sum of 14 + 4+ 35.5 = 53.5. But the found density is only one quarter of this number, viz., 13.375. It was at first imagined that this discrepancy was to be explained by abnormal expansion of the gas ; but with such a supposition, of course, Avogadro's law could not hold. Other sub stances which show the same " abnormal densities" are pentachloride of phosphorus and sulphuric acid. To ex plain this abnormality, Henri SaintClaire Deville pro pounded the idea that such substances do not go into the state of gas as compounds, but that they split into simpler components, each of which has its usual density, and a mixture of the components will exhibit a mean density. Thus, if ammonium chloride be imagined to decompose into ammonia and hydrogen chloride on changing into gas, then the density of the supposed ammonium chloride gas will be the mean of the densities of its two constituents. Ammonia has the formula NH 3 , and hydrogen chloride, HC1 ; the former has the density 8.5, and the latter, 18.25 an the mean of these two numbers is 13.375. Phosphoric chloride, which has the formula PC1 5 , splits in a similar manner into PC1 3 and C1 ; and sulphuric acid, H 2 SO 4 , into water, H 2 6, and sulphuric anhydride, SO g . To this kind of decomposition, where the bodies which are decomposed by a rise of temperature re-unite on cooling to form the origi* nal substance, Deville gave the name dissociation. It has been found possible, by taking advantage of the fact that light gases, like ammonia, pass out through an opening, or, as it is termed, " diffuse" more rapidly than heavier gases, like hydrogen chloride, to separate these gases, and thus to prove that they exist as such in the vapour of ammonium chloride ; for compounds are not decom posed into their constituents by diffusion ; hydrogen chlo ride diffuses as such, and is not split into hydrogen and chlorine.
Let us look at this dissociation from another standpoint. We know that if 2 grams of hydrogen, or 32 grams of oxygen, or 28 grams of nitrogen, or, in fact, the molecular weight of any gas expressed in grams, be caused to occupy 22,380 cubic centimeters at o C., the pressure exerted by the gas will be 76 centimeters of mercury. If the tempe rature is higher, the pressure will be increased proportionally to the increase in absolute temperature. Thus, suppose the 'temperature were 300 C., the pressure would be increased in the proportion 273 Abs. : 573Abs. :: 7 6 cms. : 1 60 cms. Now, if 53.5 grams of ammonium chloride were placed in a vacuous vessel of 22,380 cc. Capacity, and the temperature were raised to 300 C., and if no dissociation were to take place, one would expect a pressure equal to that of 1 60 cms. Of mercury. It has been found, however, that the actual pressure is twice that amount, or 320 cms. In order to account for the doubled pressure, the supposition that dissociation has taken place must again be made ; that is, in order that the pressure must be doubled, twice as many molecules must be present as one would have supposed from the weight taken. The fact of dissociation may accordingly be inferred either from a diminished density or from an increased pressure.
"Dissociation" of Salts in Solution
Few measurements of the osmotic pressure of salts have been made, owing to the difficulty in producing a membrane which shall allow water to pass, and which shall be impermeable to salts. But very numerous measurements of the depression in freezing-point and the rise in boiling-point of solutions of salts have been made ; and it has been already explained that these quantities are proportional to the osmotic pressure of the dissolved substances. It has been experimentally discovered that in all such cases the fall in freezing-point, or the rise in boiling-point is too great for the supposed molecular weight of the salt. It must be concluded that the osmotic pressure would also be increased, were it possible to measure it. But the fall in freezing-point or the rise in boiling-point does not imply a doubled osmotic pressure, when there is reason to expect it, unless the solution is very dilute. Now, if the pressure were doubled, we might argue from such cases as ammonium chloride that dissociation into two portions had occurred ; but in moderately concentrated solutions, as the pressure is not doubled, it must be concluded that the dissociation is not complete ; it is only in very dilute solutions that complete dissociation can be imagined to have taken place. Cases are known where substances in the state of gas undergo gradual dissociation, and then the pres sure does not attain its maximum until the temperature has been sufficiently raised or the pressure sufficiently reduced. The reason that this is not noticed with ammonium chloride is that the temperature of complete dissociation has been reached before the substance turns to gas.
Common salt is chloride of sodium ; its formula is NaCl ; and for long the suggestion that it dissociated into an atom of sodium and an atom of chlorine on being dissolved in water was received as too improbable to be worth consideration. There is, of course, another way out of the difficulty ; it is to suppose that a molecule of salt has the formula Na 9 Cl 2 ; in that case, 117 grams of salt (2 x 23) + (2 x 35.5) dissolved in 10,000 grams of water should produce the normal lowering of freezing point ; or, if it produced a larger lowering, it might be supposed that these complex molecules had split more or less completely into the simpler molecules, NaCl. But though the explanation suggested might account for this instance, it is incapable of accounting for the fact that chloride of barium, which is known to possess the formula BaCl 2 (or a multiple thereof), gives, in sufficiently dilute solution, a depression three times that which one would have expected from its supposed molecular weight, or that ferricyanide of potassium and ferrocyanide of potassium, the formulae of which are respectively K 8 Fe(CN) 6 and K 4 Fe(CN) 6 , should give four and five times the expected depression. But these results are quite consistent with the hypothesis that
NaCl + Aq decomposes into Na. Aq and Cl. Aq ; BaCl 2 + Aq decomposes into Ba. Aq and Cl. Aq, and Cl. Aq ; K 3 Fe(CN) 6 + Aq decomposes into K.Aq + K.Aq + K. Aq,
and Fe(CN) 6 .Aq; and K 4 Fe(CN) 6 + Aq decomposes into K.Aq + K.Aq + K.Aq
+ K.Aq, and Fe(CN) 6 .Aq.
(The symbol " Aq " stands for an indefinite but large amount of water "aqua.") Here again we are face to face with facts and an attempted explanation. The facts are that certain compounds, which have long been known as " salts," give too great a depression of the freezing-point or too great a rise of boiling-point of the solvent in which they are dissolved, corresponding to too great an osmotic pressure. It has been observed that when the dilution is sufficient the depression in each case reaches a maximum, and that that maximum is two, three, four, or five times what might be expected ; and in each case it is possible to divide the salt into two, three, four, or five imaginary por tions, which often consist of atoms, though frequently of groups of atoms.
