Читать книгу Clathrate Hydrates - Группа авторов - Страница 23

The first recorded observations of gas hydrates appear to have been made in the late 1700s. The discovery of a variety of gases in the second half of the eighteenth century made the study of gas properties a very active research topic. With researchers studying the properties of aqueous solutions of the newly discovered gases, often in unheated laboratories, it is not surprising that gas hydrates were likely made a number of times, but without a full appreciation of what was being observed. In 1786 Joseph Priestly [3], having discovered a number of “airs” (the gases NO, N₂O, HCl, NH₃, O₂, SO₂), observed that:

It is remarkable that water impregnated with vitriolic acid air (authors comment: SO₂) retains its air when it is frozen though every other kind of air (if the liquor containing it can be frozen at all) is separated from it in the act of freezing. I have now observed that this ice sinks in the liquor from which it is frozen in which it resembles the ice of oil. This is a fact which I barely mention without having any theory to account for it.

Somewhat similar observations were made while working on phlogisticated vitriolic acid (sulfurous acid) by Torbern Bergman, Professor of Chemistry at Uppsala, in his A Dissertation on Elective Attractions, translated from the Latin in 1785: [4] “This acid freezes in the same temperature as pure water; and what is remarkable, the acid elastic fluid remains in the ice, though in open vessels it forsakes the water.” From these descriptions, it is fairly clear that Bergman and Priestley had prepared SO₂ hydrate; however, as far as we know this line of investigation was neither pursued nor credited by later gas hydrate researchers.

As the general interest in the freezing point of liquids and solutions increased in the 1780s, the limits imposed by the local ambient temperature became a serious problem. Henry Cavendish, a well‐known English natural philosopher/scientist, went to some extremes to access low temperatures by sending samples across the Atlantic into Hudson Bay where experiments were carried out at some of the Hudson's Bay Company trading posts [5]. Of note are the experiments to freeze aqueous sulfuric acid solutions (H₂SO₄·2H₂O, melting point (m.p.) −37 °C, H₂SO·4H₂O, m.p. –24.5 °C), with the observations carried out in 1786 by John McNab, the Master of Henley House, on James Bay. In 1799, Cavendish became one of the founding scientists of the Royal Institution of Great Britain where he kept an active interest in Humphrey Davy's chemical experiments.

Chlorine hydrate appears to have been made [6–8] in 1785 by Bertrand Pelletier [9], a pharmacist and chemist in Paris and in 1786 by Wenceslaus Johann Gustav Karsten [10], a professor of mathematics and physics at Halle University. They observed the formation of yellow crystals when chlorine gas was cooled. There crystals were thought to be solid chlorine. During the 1790's, laboratory methods for achieving low temperatures were developed. Walker [11], an apothecary at Oxford (1790), and Lowitz [12], a chemist at St. Petersburg (1796), experimented with mixtures of salt solutions with snow and in some instances were able to freeze mercury at −40 °C. With this technology in hand, in 1799, Antoine F. de Fourcroy a lecturer in chemistry at the college of the Jardin du Roi in Paris and Louis Nicolas Vauquelin, his assistant for some time, froze a sample of liquid‐oxygenated muriatic acid gas (chlorine) to give greenish yellow crystals with a greasy texture [13]. By the time Davy came to his conclusion regarding the true nature of the material in question; the yellow solid “chlorine” was already a well‐known entity.

Humphry Davy (Figure 2.1) reported at the Bakerian Lecture [14] of the Royal Society, London, in November 1810, that:

It is generally stated in chemical books that oxymuriatic gas (authors comment: Cl₂) is capable of being condensed and crystallized at a low temperature; I have found by several experiments that this is not the case. The solution of oxymuriatic gas in water freezes more readily than pure water, but the pure gas dried by muriate of lime (authors comment: CaCl₂) undergoes no change whatever at a temperature of 40° below 0° of Fahrenheit. The mistake seems to have arisen from the exposure of the gas to cold in bottles containing moisture.

The observation that a distinct species with a characteristic high melting point formed from chlorine gas and water and the wide dissemination of this fact was enough to give Davy credit as the official discoverer of “gas hydrates.”

In 1823, Michael Faraday (Figure 2.1), working in Davy's laboratory, reported [15] the composition of chlorine hydrate as having 10 water molecules for every chlorine molecule, although he recognized the likelihood of having underestimated the chlorine content as some chlorine may have escaped the crystal during the drying process. The modern accepted value is 1.3–1.7 chlorine molecules to 10 water molecules. At Davy's suggestion, Faraday heated chlorine hydrate in a sealed tube, leading to high pressures, thus producing two immiscible liquids, water and liquid chlorine [16]. This first production of liquid chlorine led to a general method for the liquefaction of many other gases by the chemical generation of high pressure through hydrate decomposition.

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].

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: C_num – 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 water at a pressure near that required for liquefaction, followed by sudden détente of pressure to produce a crystal nucleus, and then increase of pressures to above the value at which the wall of the containing vessel becomes coated with hydrate crystals. With a subsequent reduction of pressure below the value of bulk hydrate formation, which did not depend on the relative amount of water and carbon dioxide present, the hydrate disappeared. To determine the composition, Wróblewski [45] volumetrically measured the quantity of CO₂ gas which combined with a small weighed amount of water. Accounting for non‐ideality corrections for the gas, he found the stoichiometry of CO₂·8.01H₂O as the average of 19 analyses at 16 atm and 0 °C, with a standard deviation in the hydration number of ±0.54. A further study [46] involved the role played by the abrupt fall of pressure during détente and crystallization. He promoted the principle that hydrates can only form when the concentration of dissolved gas in the aqueous solution matches its concentration in the hydrate. This condition is not normally met with carbon dioxide which becomes a liquid at a pressure well below that at which its concentration in liquid water becomes equal to its hydrate composition. He believed, however, that the cooling produced by détente could produce the requisite increase in solubility. The principle of equal concentrations was not generally true for the other known hydrates, in particular for the case of methane hydrate. That it had credibility reflects the rather poor understanding of phase equilibria at the time. Wróblewski recognized that the cooling normally produced ice as well as hydrate and insisted that all of the water would be converted to hydrate only if the relative amount of water was very small and its surface area very large.

By 1880, eight gas hydrates were known, usually observed as octahedral crystals formed at relatively low gas pressures. These had melting points above 0 °C and compositions close to eight waters per molecule of gas. The researchers involved were often on the track of other projects, so that hydrate discoveries were incidental and sustained efforts to study gas hydrates as a distinct class of materials were not made. However, after 1880, there were important changes as the tools to work at higher pressures became available, along with more reliable methods of hydrate synthesis and characterization. The laws of chemical thermodynamics were also being established and used during this time, which put the analysis of hydrate formation on a sound conceptual framework. There were concerted efforts to understand the gas hydrates as a distinct class of materials as several researchers used physical chemistry techniques to study gas hydrates for their doctoral dissertations.