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 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.
2.3 The Phase Rule
Early studies on gas hydrates had shown that the equilibrium pressure of formation of the hydrate from (or decomposition of the hydrate into) liquid water and gas depended only on temperature (i.e. the equilibrium is univariant), and this equilibrium pressure increased with increase of temperature for chlorine [43], phosphine [42], hydrogen sulfide [42], and carbon dioxide hydrates [44]. This behavior was similar to that observed when a solid decomposed into a solid and a gas and was known as Debray's law after the recent observations of the dissociation pressures of calcium carbonate and a variety of stoichiometric salt hydrates.
The proper definition of the phase relationship possible in gas–water systems capable of hydrate formation is due to H.W. Bakhuis Roozeboom (Figure 2.3), working at Leiden. In a 70‐page long memoire [47] published in 1884, Roozeboom gave the results of his careful studies of the gas hydrates of sulfur dioxide, chlorine, bromine, and hydrogen chloride; work which constituted the subject of his doctoral thesis. He confirmed the applicability of Debray's Law and showed that the aqueous solution which coexists with the hydrate at different temperatures has the same vapor tension (i.e. vapor pressure) as the hydrate and that its concentration increases with temperature as the pressure is increased. He disproved Wróblewski's contention that the equilibrium solution must have the same composition as the hydrate. He distinguished between the critical decomposition in an open vessel, where the dissociation pressure is one atmosphere, and the critical decomposition temperature in a closed vessel where the gas liquefies. Neither of these temperatures was a critical temperature in the absolute sense of a temperature above which gas hydrate can never exist. Roozeboom called the highest temperature at which SO2 hydrate can exist in a closed tube a “point de discontinuité” since the hydrate at least could continue to exist (see point Q2, Figure 2.3) in equilibrium with liquid water and liquefied gas at a higher temperature with the application of increased external pressure. In this assertion, he disagreed with Cailletet. He confirmed the observations of Cailletet and Bordet [42] and Wróblewski that, although gas hydrates were normally difficult to crystallize from solution under gas pressures well in excess of those required for hydrate stability (which with today's understanding is related to the kinetic barrier to hydrate nucleation), this difficulty was not encountered for solutions in which hydrate had been present previously.
Figure 2.3 Pioneers of clathrate science in the late 1800s and early 1900s. From left to right, Hendrik Willem Bakhuis Roozeboom, Robert Hippolyte de Forcrand, and Paul Ulrich Villard. Sources: Original photograph by Albert Greiner, reproduced with permission from the Allard Pierson Museum, University of Amsterdam, Reproduced with permission from Université de Montpellier, Reproduced courtesy of the Archives de l'Académie des Sciences, 23, Quai de Conti, 75006 Paris, France.
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