Clathrate Hydrates. Группа авторов

      Figure 2.1 Pioneers of clathrate science from the early 1800s. From left to right, Sir Humphry Davy. Stipple engraving by E. Scriven after Sir T. Lawrence, 1810/11. Credit: Wellcome Collection. Attribution 4.0 International (CC BY 4.0). Michael Faraday. Photograph by Henry Dixon & Son Ltd. Credit: Wellcome Collection. Attribution 4.0 International (CC BY 4.0). Carl Jacob Löwig in Zürich Lith. Von Orell Füssli & Cie., [zwischen 1840 und 1850?]. Zentralbibliothek Zürich, GRA 4.132. Public Domain Mark.

      Soon after, in 1828, Carl Löwig (Figure 2.1) [17], Professor of Chemistry at the University of Zürich, and one of the co‐discoverers of bromine, made orange–yellow crystals of bromine hydrate by passing bromine gas through a damp tube at 4–6°, or, by adding water to liquid bromine at 0 °C. The hydrate decomposed into bromine water and liquid bromine when heated above 15 °C [18]. Analysis of the hydrate gave 10 waters per bromine molecule, an identical composition to that found by Faraday for chlorine hydrate, but a value which was later found to be incorrect. Löwig also reported the formation of BrI and ClBr hydrates in the dissertation Das Brom und seine chemischen Verhältnisse (Bromine and its Chemical States), published in Heidelberg in 1829 [19], although later attempts to repeat the work on BrI hydrate were not successful [20].

      The following year, August A. de la Rive, a Swiss physicist [21], reported a hydrate of sulfur dioxide during attempts to liquefy SO₂. By measuring the volume of gas liberated and the weight of water left after decomposition at 4–5 °C, he found the composition to be about SO₂·14H₂O. From similarities to chlorine hydrate, he reasoned that the true composition might well be near to the 1 : 10 ratio of chlorine hydrate. From the limited amount of data available, he then reasoned that gas hydrate formation might be a property common to many gases, e.g. ammonia and H₂S, anticipating the existence of a large number of gas hydrates. However, ammonia was found to form a number of “low hydrates,” namely hemi‐, mono‐, and dihydrates [22, 23], 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.

A slightly yellow hydrate of chlorine dioxide was observed by Auguste Millon [30], professor at the Val‐de‐Grace military hospital in Paris, and again in 1869 by Moritz Brandau [31]. This was not recognized as a gas hydrate until 1906 as a result of more detailed studies by William Crowell Bray [32], a Canadian working at Leipzig, later becoming Professor of Chemistry at the University of California at Berkeley. Of note, chlorine dioxide is a free radical and readily explodes unless it is dissolved in water or trapped in the gas hydrate lattice. As such, the chlorine dioxide clathrate hydrate is an example of the safe storage of normally unstable materials in gas hydrate form, a concept recently applied to the storage of alkyl free radicals and ozone in clathrate hydrate phases.

      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:

En raison de cette même circonstance, ces corps sont éminemment propres à l'examen des modifications apportées aux propriétés physiques des deux éléments d'une combinaison par le seul fait de leur réunion. En effet, les deux composants subsistent intégralement dans le composé; ils y conservent un état moléculaire aussi voisin que possible de celui qu'ils possèdent à l'état de liberté.¹

      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].