to the discovery and characterization of weak intermolecular forces. Neither covalent nor ionic, nevertheless such forces are able to control the assembly of complex structures. The work on clathrate hydrates presaged the field of supramolecular chemistry – that is, “chemistry beyond the molecule.” Today, it is known that multiple weak interactions, e.g. van der Waals forces, hydrogen bonding, and halogen bonding, play a major role in the construction of complex materials in chemistry, biology, and materials science. This burgeoning field was duly recognized by the award of the 1987 Nobel Prize in Chemistry to Jean‐Marie Lehn, Donald J. Cram, and Charles J. Pederson.
The first ∼150 years of clathrate hydrate science is presented following the order in which these materials revealed themselves to the scientific community. In writing this chapter on the research that shaped the early scientific knowledge of clathrate hydrates, we acknowledge as major resources W. Schröder's “The History of Gas Hydrates” [1] and Donald Davidson's notes on the history of clathrate hydrates which he had started to compile into a manuscript chapter shortly before his death in 1986 [2].
2.2 The First Gas Hydrates
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