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Inflammable Ice: Clathrate Hydrates

Following the "Deepwater Horizon" oil platform explosion and oil leak in the Gulf of Mexico, attempts to pipe leaking oil to the surface were thwarted by methane hydrates that clogged the "containment domes" and pipelines.

Fig. 1. Clathrates clog containment dome

Methane hydrates are a type of "clathrate" or "cage" compound where methane (or other small nonpolar molecues) are trapped in "cages" of ice. The structure of the "clathrate", and the flammability of the solids, are illustrated below. The formation of hydrates or clathrates is made possible by the extensive hydrogen bonding in water.

Fig. 2. Clathrate structure and burning methane clathrate,[1]
Fig. 3. Structure of a methane clathrate block embedded in the upper meter of sediment off Oregon, USA at a depth of 1200 meters[2].

Clathrates form under high pressure and/or low temperatures in petroleum pipelines in the northcountry, but they are also deposits formed naturally at the ocean floor where methane is produced by microbial action. They may contain twice the amount of carbon in all other known fossil fuel reserves[3]. One litre of methane clathrate solid contains approximately 168 litres of methane gas (at STP).

The average methane clathrate hydrate composition is 1 mole of methane for every 5.75 moles of water, but it is variable, with typically 14 water molecules surrounding the methane molecule. The density is around 0.9 g/cm3, compared to 0.9167 g/cm³ for ice at 0°C, whereas water has a density of 0.9998. The similarities in density give a clue to the structure of hydrates.


The Structure of Ice

Ice is the most important example of hydrogen bonding. There is clear evidence of hydrogen bonding in the structure of the solid. The Figures show two computer-drawn diagrams of the crystal lattice of ice. In the model we can clearly see that each O atom is surrounded by four H atoms arranged tetrahedrally.

Figure 4. Two computer images of the structure of ice. The water molecules have been arranged, so that each oxygen atom is surrounded by four hydrogen atoms in tetrahedral geometry. Two of these atoms are covalently bound to oxygen, while the other two are hydrogen bonding with the oxygen.
Fig. 5. Hydrogen bonded structure of ice

The two hydrogen atoms at a distance of 99 pm from the oxygen are clearly covalently bonded to the O atom. The other two are at a distance of 177 pm, and must be hydrogen bonded to the oxygen in question (and covalently bonded in the water molecule to which they belong). If the hydrogen bond did not exist, the distance between the H of one water molecule and the oxygen of another would be closer to 260 pm, the sum of the radii of the H and O atoms.

Image:Covalent and H Bonding in Water.jpg

The tetrahedral orientation of H atoms around O atoms which results from hydrogen-bond formation has a profound effect on the properties of ice and liquid water, and provides the "cage" for clathrate formation. In the space-filling diagram of ice, most of the electron density of each H and O atom is enclosed by a boundary surface. As you can see, hydrogen bonding causes the H2O molecules to adopt a rather open structure with hexagonal channels running through it. These channels contain an almost perfect vacuum-in them there is a little electron density from the surrounding atoms, but nothing else. With small modifications, the hydrogen bonded structure is perfect for the inclusion of small molecules, as shown in Figure 2.

When ice melts, some of the hydrogen bonds are broken and the rigid crystal lattice collapses somewhat. The hexagonal channels become partially filled, and the volume of a given amount of H2O decreases. This is the reason that ice is less dense than water and will float on it. As the temperature is raised above 0°C, more hydrogen bonds are broken, more empty space becomes occupied, and the volume continues to decrease. By the time 4°C has been reached, increased molecular velocities allow each H2O molecule to push its neighbors farther away. This counteracts the effect of breaking hydrogen bonds, and the volume of a given amount of H2O begins to increase with temperature.

Effects of Expansion During Freezing

Most solids expand when they melt, and the corresponding liquids expand continually with increasing temperature, so the behavior of water is rather unusual. It is also extremely important in the environment. When water freezes in small cracks in a rock, the greater volume of the ice can split the rock into smaller pieces. These eventually become able to support plant life, and so water contributes to the formation of fertile soil. The same process happens to roadways, and is the reason for new cracks and potholes seen on roads after a cold winter. The ice bomb experiment, seen below, is perhaps the the most dramatic example of water expanding when frozen.

In the video, water is poured into a cast iron container, which is tightly sealed. The container is then placed in a acetone/dry ice sludge, which is at a temperature of 77°C. After a short period of time, the ice freezes, expands, and causes the cast iron container to explode, blasting off the cover of the acetone/dry ice bath, and spraying the bath itself everywhere. Even though the cast iron container had ⅛ inch thick sides, the pressure of the expanding ice was still able to blow it apart.

Since water has maximum density at 4°C, water at that temperature sinks to the bottom of a deep lake, providing a relatively uniform environment all year around. If ice sank to the bottom, as most freezing liquids would, the surface of a lake would not be insulated from cold winter air. The remaining water would crystallize much more rapidly than it actually does. In a world where ice was denser than water, fish and other aquatic organisms would have to be able to withstand freezing for long periods.

Heat Evidence for Hydrogen Bonding

Hydrogen bonding also contributes to the abnormally large quantities of heat that are required to melt, boil, or raise the temperature of a given quantity of water. Heat energy is required to break hydrogen bonds as well as to make water molecules move faster, and so a given quantity of heat raises the temperature of a gram of water less than for almost any other liquid. Even at 100°C there are still a great many unbroken hydrogen bonds, and almost 4 times as much heat is required to vaporize a mole of water than would be expected if there were no hydrogen bonding. This extra-large energy requirement is the reason that water has a higher boiling point than any of the other hydrides.


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