ТОП просматриваемых книг сайта:
Clathrate Hydrates. Группа авторов
Читать онлайн.Название Clathrate Hydrates
Год выпуска 0
isbn 9783527695065
Автор произведения Группа авторов
Издательство John Wiley & Sons Limited
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].
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.
Cailletet and Bordet measured the pressures of formation of H2S and PH3 hydrates over a range of temperatures and found each formation temperature to correspond to a unique value of pressure. F. Isambert had already shown [43] the univariant nature of the equilibrium pressure of chlorine hydrate by measuring the relative amounts of water and gas present in the hydrate within wide temperature and pressure limits (see below for further discussion). Cailletet referred to the upper temperature at which the gas hydrate dissociation curve ended upon liquefaction of the gas (28 °C for PH3 hydrate and 29 °C for H2S hydrate) as a critical temperature, above which hydrate cannot be formed under any pressure. In modern language, this critical temperature identifies the upper quadruple point (Q2, hydrate–aqueous solution–gas–liquid phases) of the phase diagram, see Section 2.3. The existence of a critical temperature at Q2 is not strictly true of all gas hydrates, but it is a reasonable approximation.
L.P. Cailletet and L. Bordet [42] found that the hydrate obtained from gas consisting of equal volumes of PH3 and CO2 was not simply a mixture of the pure hydrate of CO2 and a pure hydrate of PH3. The hydrate formed by the mixed gases had a critical temperature of 20 °C, while the critical temperature of the pure PH3 hydrate was 28 °C and that of CO2 hydrate was below 7 °C. This result, although it describes a feature unique to gas hydrates, was not further discussed in terms of the formation of mixed hydrates.
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