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 stabilize known hydrates [74]. Starting with initial success with acetylene, he also found that ethylene, carbon dioxide, and SO2 could perform that function.
In the new century, de Forcrand initiated a new approach to the determination of hydrate compositions in recognition of the fact that direct determinations were difficult and prone to errors. He generalized Trouton's rule, proposed in 1887, that the entropy of vaporization for various kinds of liquids at their boiling points is almost the same value, about 85–88 J K−1 mol−1 [75]. The entropy of vaporization is defined as the ratio between the enthalpy of vaporization and the boiling temperature. de Forcrand calculated compositions for all of the known hydrates, first improving doubtful data when necessary. The results of his calculations are shown in Table 2.2 [28], of which about half of the entries appear to support Villard's rule. Except for outliers Ar and Br2, for the other entries, both the heats of dissociation to form ice Q(ice) and water Q(water), and the hydration number generally increased with molecular weight to give up to eight waters/guest.
Further progress in determining hydrate compositions virtually ground to a halt, as neither direct, nor indirect, methods were able to give a convincing explanation of the apparent complexity of the variable hydrate compositions. For instance, de Forcrand, from previously obtained data, calculated the composition of the chloroform–H2S hydrate to be CHCl3·2H2S·19H2O. This composition was explained by de Forcrand in terms of the formula (CH3Cl·7H2O) + 2(H2S·6H2O); thus, two hydrates present in a 1 : 2 ratio. This formula was then taken to be common to all sulfhydrated hydrates (binary hydrates with H2S). The composition found that for chloroform hydrate, CHCl3·18H2O in 1885 by Chancel and Parmentier [61] was ascribed to the presence of a large excess of water, although CHCl3·17H2O is the true formula.
More complexities arose from de Forcrand's efforts to investigate hydrate formation by the noble gases [76], in particular argon hydrate after it having been reported by Villard [28, 77]. He was able to make krypton hydrate, and from the dissociation behavior and heats of formation, he arrived at a composition of Kr·5.08H2O, and a redetermination of the value for Ar hydrate led to a composition of Ar·5.5H2O. Eventually, he was able to form Xe hydrate and determined its composition to be Xe·6.6H2O [78]. Rounding off, Ar and Kr then have a hydration number of 5 or 6; however, xenon's value then would be 6 or 7, which again led to speculation why these rather similar noble gases would have different hydration numbers. There were further efforts made to confirm or refute Villard's rule, but without much success either way. The formation of hydrates of noble gas indicated that the chemists of the day realized that the water–gas interactions in hydrates were not chemical in nature.
Table 2.2 De Forcrand's hydrate compositions obtained using calorimetric data [1, 28].
| Guest | Tboiling (K) | Tdissoc. (K) | Q(ice) (cal) | Q(water) (cal) |
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