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Clathrate Hydrates


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all of which melt below the melting point of NH3.

      Hydrogen sulfide indeed did fit into de la Rive's scheme as H2S hydrate was made in 1840 by the well‐known chemist Friedrich Wöhler [24], Professor of Chemistry and Pharmacy at Göttingen and Inspector General of all apothecaries in Hannover. He prepared the hydrate at ordinary temperatures by allowing hydrogen sulfide to react with water under some pressure in a sealed tube, and alternatively, at ambient pressure by passing H2S gas into a water–alcohol mixture at −18 °C. As was by now not unusual, the hydrate composition was not easily established. Over the years, the hydration numbers for H2S hydrate were reported as 15 [25], 12 [26], 7 [27], and 6 [28]. A value of 5 was also reported [29], based on 28 measurements that lay within 0.2 of 5.3. These numbers illustrate a universal custom that persisted until the 1950s, that is, rounding off the results of hydration number measurements so as to report these as simple integral proportions, following the usual procedures in inorganic chemistry and the law of definite proportions since the time of Proust and Dalton.

      Hydrates of organic molecules were first synthesized by Marcellin Berthelot [33], who reported the hydrates of methyl bromide and methyl chloride in 1856. Although relative latecomers to the gas hydrate family, hydrates of organic molecules are now in the great majority of hydrates made and studied. The future Nobel laureate, J.W.F. Adolf von Baeyer, independently prepared methyl chloride hydrate simply by cooling a saturated aqueous solution to 6 °C [34]. Berthelot, like many workers after him, was struck by the weakness of the forces which hold the components of gas hydrate together:

      Berthelot also identified as CS2 hydrate, the unstable “snow” that formed when a CS2 solution was filtered in an air stream. This sparked a debate, lasting 40‐years, among workers in more than a dozen laboratories, about the existence and properties of this kind of “hydrate.” The current view is that CS2 by itself does not form a gas hydrate, although it will do so with a “help‐gas.”

      The reactions of ethylene oxide were first examined in 1863 by Charles‐Adolphe Wurtz [35], who reported the formation of ethylene oxide hydrate but whose composition remained undetermined. In 1922, melting point–composition diagrams of ethylene oxide hydrate gave a hydration number between 5 and 8 and a congruent melting point of 11 °C [36, 37]. As a completely water‐soluble material, ethylene oxide differed from previous hydrate formers that were nearly insoluble or only weakly soluble in water. Finally, it was X‐ray diffraction that showed ethylene oxide could indeed form a gas hydrate isostructural with other known gas hydrates [38]. Today, a variety of water‐soluble materials are known as hydrate formers, encompassing ethers, ketones, aldehydes, alcohols, and others.

      In 1878, Louis Paul Cailletet, a physicist and master in iron‐working, published [39] a description of the elegant apparatus and techniques ultimately used by him to generate high pressures and low temperatures necessary to liquefy even the “permanent” gases, Figure 2.2. The Cailletet apparatus was widely copied and modified for the use in many kinds of experiments requiring high pressures. This included the preparation of new gas hydrates and the determination of pressure–temperature regimes under which gas hydrates were stable. The cooling effect frequently produced by sudden reduction of pressure on a gas was apparently first noticed by Cailletet and later was put to good use in inducing the crystallization of gas hydrates from metastable or supersaturated solutions. The first gas hydrate to be produced using this method was acetylene hydrate, which resulted from cooling a mixture of liquid acetylene, linseed oil, and water to 0 °C [41]. A few years later, the gas hydrate of phosphine was prepared by subjecting a mixture of liquid PH3 and water to compression, cooling, and a sudden release of pressure causing an abrupt fall in temperature, a procedure known as détente [42].

Photograph depicts Cailletet apparatus [40] showing the hand-driven hydraulic pumps (M and O, lying horizontally on the table) used to compress (via inert mercury) the cooled sample, which lies in the vertical cylinder (m), near the center of the figure. Crystal formation can be seen through the glass sample holder. After compression of the sample and cooling through an outer jacket, the hydraulic pressure on the sample is released, leading to expansion of the sample and a further temperature drop (détente). The pressure on the sample is measured with the mercury manometer on the right (N and N′).

      Figure 2.2 Cailletet apparatus [40] showing the hand‐driven hydraulic pumps (M and O, lying horizontally on the table) used to compress (via inert mercury) the cooled sample, which lies in the vertical cylinder (m), near the center of the figure. Crystal formation can be seen through the glass sample holder. After compression of the sample and cooling through an outer jacket, the hydraulic pressure on the sample is released, leading to expansion of the sample and a further temperature drop (détente). The pressure on the sample is measured with the mercury manometer on the right (N and N′). Source: Tissandier [40], reproduced with permission from: Cnum – Conservatoire numérique des Arts et Métiers.

      Zygmunt Florenty Wróblewski, while working in Jules Henri Debray's laboratory at the École Normale Supérieure in Paris, made use of the Cailletet apparatus in the studies of the solubility of carbon dioxide which led to the discovery of carbon dioxide hydrate in 1882 [44]. Since CO2 hydrate requires a somewhat higher pressure for stability (e.g. 12.3 atm at 0 °C) than hydrates prepared previously, it was prepared by compressing CO2 gas over