Part 12 (1/2)

[Ill.u.s.tration: FIG. 34.]

The vapour pressure of the different systems of sodium sulphate and water can best be studied with the help of the diagram in Fig. 34.[216] The curve ABCD represents the vapour-pressure curve of the saturated solution of anhydrous sodium sulphate. GC is the pressure curve of decahydrate + anhydrous salt, which, as we have seen, cuts the curve ABCD at the transition temperature, 32.6. Since at this point the solution is saturated with respect to both the anhydrous salt and the decahydrate, the vapour-pressure curve of the saturated solution of the latter must also pa.s.s through the point C.[217] As at temperatures below this point the solubility of the decahydrate is less than that of the anhydrous salt, the vapour pressure of the solution will, in accordance with Babo's law (p. 126), be higher than that of the solution of the anhydrous salt; which was also found experimentally to be the case (curve HC).

{141}

In connection with the vapour pressure of the saturated solutions of the anhydrous salt and the decahydrate, attention must be drawn to a conspicuous deviation from what was found to hold in the case of one-component systems in which a vapour phase was present (p. 31). There, it was seen that the vapour pressure of the more stable system was always _lower_ than that of the less stable; in the present case, however, we find that this is no longer so. We have already learned that at temperatures below 32.5 the system decahydrate--solution--vapour is more stable than the system anhydrous salt--solution--vapour; but the vapour pressure of the latter system is, as has just been stated, lower than that of the former.

At temperatures above the transition point the vapour pressure of the saturated solution of the decahydrate will be lower than that of the saturated solution of the anhydrous salt.

This behaviour depends on the fact that the less stable form is the more soluble, and that the diminution of the vapour pressure increases with the amount of salt dissolved.

With regard to sodium sulphate heptahydrate the same considerations will hold as in the case of the decahydrate. Since at 24 the four phases heptahydrate, anhydrous salt, solution, vapour can coexist, the vapour-pressure curves of the systems hydrate--anhydrous salt--vapour (curve EB) and hydrate--solution--vapour (curve FB) must cut the pressure curve of the saturated solution of the anhydrous salt at the above temperature, as represented in Fig. 34 by the point B. This const.i.tutes, therefore, a second quadruple point, which is, however, metastable.

From the diagram it is also evident that the dissociation pressure of the heptahydrate is higher than that of the decahydrate, although it contains less water of crystallization. The system heptahydrate--anhydrous salt--vapour must be metastable with respect to the system decahydrate--anhydrous salt--vapour, and will pa.s.s into the latter.[218]

Whether or not there is a temperature at which the vapour-pressure curves of the two systems intersect, and below which the heptahydrate becomes the more stable form, is not known.

{142}

In the case of sodium sulphate there is only one stable hydrate. Other salts are known which exhibit a similar behaviour; and we shall therefore expect that the solubility relations.h.i.+ps will be represented by a diagram similar to that for sodium sulphate. A considerable number of such cases have, indeed, been found,[219] and in some cases there is more than one metastable hydrate. This is found, for example, in the case of nickel iodate,[220] the solubility curves for which are given in Fig. 35. As can be seen from the figure, suspended transformation occurs, the solubility curves having in some cases been followed to a considerable distance beyond the transition point. One of the most brilliant examples, however, of suspended transformation in the case of salt hydrates, and the sluggish transition from the less stable to the more stable form, is found in the case of the hydrates of calcium chromate.[221]

[Ill.u.s.tration: FIG. 35.]

In the preceding cases, the dissociation-pressure curve of the hydrated salt cuts the vapour-pressure curve of the saturated {143} solution of the anhydrous salt. It can, however, happen that the dissociation-pressure curve of one hydrate cuts the solubility curve, not of the anhydrous salt, but of a lower hydrate; in this case there will be more than one stable hydrate, each having a stable solubility curve; and these curves will intersect at the temperature of the transition point. Various examples of this behaviour are known, and we choose for ill.u.s.tration the solubility relations.h.i.+ps of barium acetate and its hydrates[222] (Fig. 36).

[Ill.u.s.tration: FIG. 36.]

At temperatures above 0, barium acetate can form two stable hydrates, a trihydrate and a monohydrate. The solubility of the trihydrate increases very rapidly with rise of temperature, and has been determined up to 26.1.

At temperatures above 24.7, however, the trihydrate is metastable with respect to the monohydrate; for at this temperature the solubility curve of the latter hydrate cuts that of the former. This is, therefore, the transition temperature for the trihydrate and monohydrate. The solubility curve of the monohydrate succeeds that of the trihydrate, and exhibits a conspicuous point of minimum solubility at about 30. Below 24.7 the {144} monohydrate is the less stable hydrate, but its solubility has been determined to a temperature of 22. At 41 the solubility curve of the monohydrate intersects that of the anhydrous salt, and this is therefore the transition temperature for the monohydrate and anhydrous salt. Above this temperature the anhydrous salt is the stable solid phase. Its solubility curve also pa.s.ses through a minimum.

The diagram of solubilities of barium acetate not only ill.u.s.trates the way in which the solubility curves of the different stable hydrates of a salt succeed one another, but it has also an interest and importance from another point of view. In Fig. 36 there is also shown a faintly drawn curve which is continuous throughout its whole course. This curve represents the solubility of barium acetate as determined by Krasnicki.[223] Since, however, three different solid phases can exist under the conditions of experiment, it is evident, from what has already been stated (p. 111), that the different equilibria between barium acetate and water could not be represented by one _continuous_ curve.

Another point which these experiments ill.u.s.trate and which it is of the highest importance to bear in mind is, that in making determinations of the solubility of salts which are capable of forming hydrates, it is not only necessary to determine the composition of the solution, but _it is of equal importance to determine the composition of the solid phase in contact with it_. In view of the fact, also, that the solution equilibrium is in many cases established with comparative slowness, it is necessary to confirm the point of equilibrium, either by approaching it from higher as well as from lower temperatures, or by actually determining the rate with which the condition of equilibrium is attained. This can be accomplished by actual weighing of the dissolved salt or by determinations of the density of the solution, as well as by other methods.

{145}

2. _The Compounds formed have a Definite Melting Point._

In the cases which have just been considered we saw that the salt hydrates on being heated did not undergo complete fusion, but that a solid was deposited consisting of a lower hydrate or of the anhydrous salt. It has, however, been long known that certain crystalline salt hydrates (_e.g._ sodium thiosulphate, Na_{2}S_{2}O_{3},5H_{2}O, sodium acetate, NaC_{2}H_{3}O_{2},3H_{2}O) melt completely in their water of crystallization, and yield a liquid of the _same composition_ as the crystalline salt. In the case of sodium thiosulphate pentahydrate the temperature of liquefaction is 56; in the case of sodium acetate trihydrate, 58. These two salts, therefore, have a definite melting point.

For the purpose of studying the behaviour of such salt hydrates, we shall choose not the cases which have just been mentioned, but two others which have been more fully studied, viz. the hydrates of calcium chloride and of ferric chloride.

Solubility Curve of Calcium Chloride Hexahydrate.[224]--Although calcium chloride forms several hydrates, each of which possesses its own solubility, it is nevertheless the solubility curve of the hexahydrate which will chiefly interest us at present, and we shall therefore first discuss that curve by itself.

[Ill.u.s.tration: FIG. 37.]

The solubility of this salt has been determined from the cryohydric point, which lies at about -55, up to the melting point of the salt.[225] The solubility increases with rise of temperature, as is shown by the figures in the following table, and by the (diagrammatic) curve AB in Fig. 37. In the table, the numbers under the heading ”solubility” denote the number of grams of CaCl_{2} dissolved in 100 grams {146} of water; those under the heading ”composition,” the number of gram-molecules of water in the solution to one gram-molecule of CaCl_{2}.

SOLUBILITY OF CALCIUM CHLORIDE HEXAHYDRATE.

----------------------------------------- Temperature.

Solubility.

Composition.

----------------------------------------- -55

42.5

14.5 -25

50.0

12.3 -10

55.0

11.2 0

59.5

10.37 10

65.0

9.49 20

74.5