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 over water in a sealed tube, he discovered that complete transformation of water to hydrate could take weeks. He developed glass apparatus that could be charged with a weighed amount of water to which mercury was added, which when shaken, would provide mixing. The water was then frozen to ice at −20 °C, at which point the air in the tube was replaced with the hydrate forming gas. The tube was then sealed with wax, and the ice was allowed to melt, forming a hydrate “mash” on the tube walls, see Figure 2.6, with all of the water converted to hydrate. Cooling the tube to −20 °C and melting the wax allowed the excess gas to escape, meanwhile keeping the hydrate intact. Finally, the hydrate was decomposed by warming, and the amount of released gas was measured volumetrically. The procedure was further refined to allow work under pressure at 0 °C. In addition to composition studies, Villard added microscopic observation of the hydrate crystals to obtain their morphology and their interaction with polarized light as well as calorimetric measurements of the heat of decomposition. Based on the regularity of the results obtained for hydrates formed by a number of gaseous hydrate formers, he proposed the following definition of gas hydrates: [72]
Figure 2.6 Villard's apparatus for hydrate formation and characterization. On the left‐hand side, the liquid mercury is at the bottom of the tube, with liquid water above it, and gas in the container on top of the water (labeled 1). The tip of the tube, a, is sealed with wax. The apparatus is inverted (right‐hand side) and shaken with the mercury agitating the water and the gas to form hydrate. After hydrate formation, followed by decomposition, the released gas goes through compartment 2 to a gas measuring device. Source: Adapted from Villard [68], reproduced with permission from the Bibliothéque National de France.
Les combinaisons dissociables, susceptibles d'exister seulement à l'état solide, formées par l'eau avec divers gaz, sont isomorphes entre elles, cristallisent dans le système cubique, et leur constitution est exprimée par formule générale M, 6H2O, M représentant une molécule du gaz considéré.8
The liquid hydrates were treated as being similar to the gas hydrates except that their decomposition temperatures were all close to 0 °C, thus distinguishing them from the gas hydrates which were stable to higher temperatures. It was noted that under a pressure of a “helper” gas, the liquid hydrate decomposition temperatures increased markedly, with the decomposition of ethyl chloride hydrate rising from 4.8 to 5.5 °C under 23 atm of hydrogen or 2.5 atm of oxygen, as mentioned above. Mistakenly, Villard asserted that the “helper” gases did not participate in hydrate formation and proposed a thermodynamically untenable explanation for these results.
It did not take long for “Villard's rule” regarding the M·6H2O composition of the hydrate phases to be challenged, as it is clear that the direct compositional analysis of gas hydrate depends on the purity of the hydrate material prepared. Values deviating from Villard's rule easily could be attributed to either excess water or excess hydrate former associated with or trapped in the solid hydrate. W. Hempel and J. Seidel's [73] experiment on the determination of the composition of CO2 hydrate is worthy of note. CO2 hydrate was prepared by sealing water and “carbonic acid” (CO2) in a sealed tube at −79 °C, and allowing the tube to warm to room temperature. When the two liquid layers that formed were cooled to 0 °C, the hydrate formed readily. The sealed tube was again cooled to −79 °C, the tube broken open and fitted to a capillary delivery pipe, and the contents were allowed to warm slowly. The evolved CO2 gas was collected in a gasometer over mercury. After the non‐bound CO2 escaped, gas evolution ceased almost completely at −25 °C only to start again at −2 °C. Vigorous effervescence was then evident again between 0 and 15 °C. The hydration number derived from the known amount of water and the evolved gas depended on whether all of the gas released between −25 and –2 °C was attributed to hydrate decomposition. Thus, the experiment was not conclusive in testing Villard's rule. However, the slow release of gas above −25 °C may well be attributed to a “self‐preservation” effect, as the pressure over CO2 hydrate reaches 1 atm at −55 °C, and it would have been expected that most of the hydrate would have decomposed well below −2 °C. The self‐preservation effect of hydrates has been active topic of research and is discussed in Chapter 13.
In 1897, de Forcrand and Thomas initiated new studies on double hydrates to see if other help gases in addition to H2S and H2Se could be found that might
