The boiling points of the substance are those reported by de Forcrand. Other small halogenated alkanes which did not form clathrate hydrates were listed by de Forcrand.
II Composition determined by de Forcrand by two component analysis.
III Composition determined by de Forcrand by three component analysis.
H Likely structure H hydrate formers synthesized by de Forcrand.
Source: Adapted from Schröder [1], de Forcrand [26].
As evidence for the close similarity of many of the double hydrates with hydrogen sulfide, de Forcrand found that chemical analysis for all three components of the nine hydrates marked “III” in Table 2.1 and for two components (the third being determined by difference) of the seven hydrates marked “II” all gave the same composition, namely M·2H2S·23H2O. In hindsight, the compounds listed in Table 2.1 probably represent two different hydrate structures, most of them belonging to the structure II (sII) hydrate family, the compounds flagged with “H” likely are hexagonal structure H (sH or HS‐III) hydrate formers, see Chapter 3 for further discussion.
Another common feature of many of the double hydrates was the morphology of their crystals. The chloroform‐hydrogen sulfide hydrate was found to sublime as well‐defined crystals on the inner surface of the sealed tube. de Forcrand observed of this hydrate [26],
J'ai pu remarquer dans quelques cas des octaèdres presque parfaits. D'ailleurs l'examen au microscope polarisant ne peut laisser aucun doute sut la forme cubique de ces cristaux, qui n'agissent pas sur la lumière polarisée 5
Similar observations of several other of the most stable double hydrates showed cubic, cubo‐octahedral, and truncated octahedral forms (Figure 2.5a,b) which had no effect on polarized light when examined with the polarized light microscope. de Forcrand appears to have been the first to argue that the gas hydrate crystals examined belonged to the cubic system and thus were distinguishable from hexagonal ice crystals. The dissociation pressures of nine double hydrates measured by de Forcrand in the presence of liquid water and liquid hydrate former are reproduced in Figure 2.5c. These dissociation pressures and compositions of the gas phase were found to be dependent only on temperature and to be independent of the overall relative amounts of the three components. That these observations were a consequence of the thermodynamic relationship for four‐phase equilibria was not realized until Roozeboom's application of the phase rule a few years later. Figure 4 of Ref. [26], reproduced in Figure 2.5c, shows that the nature of the organic component affects the stability of the double hydrate: the dissociation pressure (mainly contributed by H2S) is 1 atm at 3 °C, for CH3CHBr2. However, the dissociation pressure is reached at a higher temperature of 18 °C for the much more stable double hydrate of CCl4.
In a first application of calorimetry to gas hydrates, de Forcrand attempted to measure the heat of dissociation of the chloroform‐hydrogen sulfide hydrate. He found that 47 cal g−1 was absorbed in the decomposition of the hydrate into the two liquids and gaseous H2S. A similar result was found for the double hydrate with ethyl bromide. He concluded that [26],
…les nombres qui en résultant prouvent que la quantité de chaleur produite est assez considérable, mais qu'elle est due surtout au changement d'état de l'eau qui entre dans le composé.6
This was a reasonable conclusion since the heat of fusion of ice was 80 cal g−1 and according to de Forcrand's composition, water made up about 70% of the mass of these hydrates. The numbers confirmed the general impression that gas hydrate formation was more akin to a freezing process than to a chemical reaction.
Except for bromine, the evidence that simple hydrates could be formed by some substances which are liquid at room temperature, mainly lay in the formation of low temperature frosts, and remained unconvincing until the characterization [61] of a hydrate of chloroform in 1885. Chancel and Parmentier [61] showed this hydrate to decompose above 0 °C and found its composition to be CH3Cl·18H2O. Chloroform hydrate was to be recognized as the first of the so‐called “liquid hydrates” which, despite their higher water content possessed many of the same properties as “gas hydrates.”
Figure 2.5 De Forcrand's results showing (a) octahedral and related crystalline forms for the binary clathrate hydrate of H2S and carbon tetrachloride; (b) the modified octahedral crystal of binary hydrate of H2S and isopropyl bromide; (c) the dissociation pressures (mmHg) as a function of temperature (degrees Celsius) for nine double hydrates in the presence of liquid water and hydrate former. Source: adapted from: Ref [26], reproduced with permission from the Bibliothéque National de France.
Formally, however, the first liquid hydrate to be reported was that of ethanethiol (ethyl mercaptan). In 1872, Hermann Müller [62] reported that in the distillation of this mercaptan from concentrated aqueous solution of the potassium salt of ethylsulfuric acid and sodium hydrosulfide, the cooling condenser became filled with the mercaptan hydrate which melted at 12 °C to give two liquid layers. From the volume of the two layers, the composition of C2H5SH·24H2O was estimated for the hydrate. Shortly thereafter, similar behavior was noted [63] by Peter Clässon, working at the University of Lund, who used elemental analysis to find the composition C2H5SH·18H2O. Clässon recognized the presence of hydrogen sulfide as an impurity in the mercaptan and attempted to remove it. Nevertheless, since ethanethiol hydrate is now known to decompose below 4 °C, it appears likely that both Müller's and Clässon's hydrates were stabilized by the presence of H2S. In 1887, Peter Klason reported [64] the formation of methyl mercaptan gas hydrate which decomposed at a temperature far higher than the boiling point of the mercaptan (12 °C).
In 1888, de Forcrand presented a series of short papers on gas hydrates which resulted from collaboration with Paul Villard (Figure 2.3). Villard had graduated from the École Normale in Paris with a teaching certificate in 1884 and became a secondary school teacher in the provinces. It appears that he started to study gas hydrates with de Forcrand in Montpellier in 1887, the year in which the latter became professor of chemistry there. In new measurements of the dissociation pressures of hydrogen sulfide hydrate [65], de Forcrand and Villard closely confirmed the earlier results of de Forcrand [25], in contrast with the higher values since measured by Cailletet and Bordet [42]. They also observed that at low temperatures:
…la présence de la glace amène des perturbations dont il est difficile de tenir compte, et qui nous ont empêché jusqu'ici d'obtenir des résultat concordants.7
Roozeboom responded [56] that he had already solved this problem [50] by showing the presence of an ice‐hydrate‐gas equilibrium which differed from the liquid‐hydrate‐gas equilibrium. Roozeboom also pointed out that the discussion by de Forcrand and Villard of
