Clathrate Hydrates. Группа авторов

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 Wróblewski's solubility principle for hydrogen sulfide and methyl chloride hydrates merely confirmed what he had himself found for other hydrates [56]; besides, thermodynamics placed no restrictions on the dissolved gas content of the liquid water phase in equilibrium with hydrate. In subsequent studies of the pressure along the methyl chloride hydrate–liquid water–gas equilibrium line [66], the compositions of the hydrates of hydrogen sulfide and methyl chloride (H₂S·7H₂O and CH₃Cl·9H₂O) were found. At the time, Roozeboom's comments were not acknowledged by other researchers.

Except for this brief period of collaboration with de Forcrand, Villard worked entirely by himself. He applied to Debray, shortly before the latter's death in 1888, for permission to work in the chemistry laboratory at his old school, l'École Normale Supérieure, where Cailletet and Wróblewski had done their work on gas hydrates. Within a matter of months, he reported [67] the formation of new gas hydrates by methane, ethane, ethylene, nitrous oxide, and acetylene, having overlooked Cailletet's prior discovery of acetylene [41] hydrate. The hydrate of methane [68] was prepared by compressing methane to about 75 atm over water near 0 °C in a Cailletet tube, followed by détente and recompression. This hydrate and the hydrate of ethylene [68] were particularly interesting since each existed above the critical point of the gas and therefore exhibited a dissociation pressure curve which was not limited at high temperatures by liquefaction of the gas. Nevertheless, from the steep rise in pressure which he observed at relatively high temperatures, Villard wrongly inferred that methane hydrate had a critical temperature of about 21.5 °C at 300 atm and ethylene hydrate had a critical point of about 18.7 °C above 60 atm. In 1890, Villard added the next higher homologue, propane to the list of hydrate formers [69].

      In a paper defining [70] the dissociation pressures of new hydrates of methyl and ethyl fluoride, Villard also observed that the hydrate formed when a cold mixture of ethyl chloride or methyl iodide with water was nucleated with ice did not survive heating to much above 0 °C. When prepared in the presence of air, however, these hydrates could be heated to about 5 °C before decomposition occurred. The stabilizing effect of air was later to become a common observation for many gas hydrates and is still a point of interest/caution in hydrate syntheses in open vessels. In testing other “help‐gases,” Villard reported that the decomposition temperature of ethyl chloride hydrate was increased from 4.8 to 5.5 °C under 23 atm of hydrogen or 2.5 atm of oxygen.

      Villard next recorded [69] the formation of hydrates by carbon tetrafluoride, fluoroform, methylene fluoride, and ethylene tetrafluoride. Like the methyl and ethyl fluoride previously used [70], these fluorocarbons were synthesized by Villard himself. In the case of carbon tetrafluoride, at least, the purity was highly questionable, as Villard found this hydrate to be stable without applied pressure at 0 °C; however, much later measurements showed that the hydrate requires 40 atm at 0 °C for stability.

      With the large number of new gas hydrates available, Villard was encouraged to attempt the measurement of the compositions of these new materials, the first two of which were N₂O and CO₂ [71]. By carefully monitoring the gas pressure of N₂O