In the Laboratory
Gas Clathrate Hydrates Experiment for High School Projects and Undergraduate Laboratories
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Melissa R. Prado, Annie Pham, Robert E. Ferazzi, Kimberly Edwards, and Kenneth C. Janda* Department of Chemistry, University of California, Irvine, Irvine, CA 92697; *
[email protected] Gas clathrate hydrates are crystalline solids with structures consisting of a lattice of hydrogen-bonded water molecules that encage small-diameter molecules of gases (1). This description alone presents a fascinating idea: a hydrophobic center in a water lattice. Gas hydrates have been the subject of research for nearly two centuries, starting with the initial detection of the chlorine hydrate in 1810 by Sir Humphrey Davy (2). Later, in the 1930s, hydrates gained further attention by blocking oil and gas pipelines (1, 3). Recently, research on these species has intensified owing to their ability to store methane and hydrogen, including large deposits of methane in ocean sediments (4). In addition to being important for practical reasons, gas clathrates present interesting problems in bonding, kinetics, and thermodynamics that have made them the subject of intense study. The combination of interesting applications and interesting science prompted us to develop an experiment to study propane clathrate hydrate suitable for undergraduate laboratories to introduce students to this fascinating solid. The experiments described here have been tested at UCI both in our undergraduate chemistry laboratory and by high school teachers during summer institutes. The project has proved to be thought provoking. Students are especially interested in seeing the clathrate hydrate, which looks similar to packed snow, burn. This is shown in Figure 1. Gas hydrates and pure water ice, in addition to their physical resemblance, are each composed of a water latticework held together and stabilized by hydrogen bonds. However, hydrates are much more complex than ice in that they
contain “host” cages where “guest” molecules are trapped by van der Waals forces (5). The water molecules in clathrate structures maintain the four hydrogen bonds to each water molecule. However, the hydrogen-bond angles are distorted from their optimum values, so that the empty clathrate lattice is less stable than water ice (6). The van der Waals interactions between the trapped gas molecules and the water molecules compensate for the less-stable water lattice. Therefore, the melting point of a gas clathrate is sensitive to gas pressure. In some cases, such as the propane clathrate discussed here, the melting point is higher than 0 ⬚C at moderate pressures. Thus we have the unusual case that the solid formed from water and a non-miscible gas has a melting point (or a dissociation temperature) higher than that of pure ice (2). The experiments described here introduce students to these fascinating crystals by synthesizing propane clathrate, which may then be burned as a demonstration (Figure 1), and whose thermodynamic and stoichiometric properties can be studied using a simple apparatus. The fact that the data are not simple to analyze forces students to think about relationships between structure, stability, and how melting points depend on pressure. This provides an opportunity for students to predict and compare how laboratory data relate to the phase diagram and to use the ideal gas law in stoichiometry calculations. Owing to its large unit cell, calculating the expected stoichiometry of propane clathrate is a challenge. Especially careful students will be able to measure the heat of vaporization from the dependence of vapor pressure on temperature and compare it to that of ice. Hazards This experiment involves propane, a flammable gas, under pressures up to 100 psi. Appropriate safety procedures are described throughout the Supplementary Material.W Unusual Materials needed The experiment requires construction of custom gas pressure cells, a coffee grinder, and access to a freezer in addition to standard laboratory equipment. The sample cells are robust and can provide years of service to offset the initial cost. We suggest using Vernier LoggerPro electronics, which are readily available and not too expensive, for data recording. A complete description and drawing of the equipment is given in the Supplementary Material.W Experimental Procedure
Figure 1. Photograph of a plug of “burning snow”, propane clathrate hydrate, prepared as described in the text.
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The procedures employed in this experiment have been modified from those of Stern et al. (7) to be practical for a
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In the Laboratory
student laboratory. Snow (powdered ice) is used as a starting material to maximize the surface area of ice exposed to propane gas. The snow is made by grinding small ice pellets into a fine powder in a precooled coffee grinder. This snow is then placed in the sample cell and exposed to the propane gas. To facilitate the conversion to propane clathrate the sample cell is stored on ice overnight. Concurrently with making the propane clathrate sample, a control sample of powdered ice with no propane is prepared. Each sample cell contains temperature and pressure probes interfaced with a computer. Some of the samples can be simply opened and burned to intrigue the students. Data collection involves monitoring the change in pressure and temperature versus time to observe both the extent of propane uptake by the ice and the extent of clathrate dissociation as the sample is warmed up. Sample Results and Discussion Points
Summary
Sample data for the pure ice sample and for the propane clathrate are shown in Figure 2. The pure ice gradually warms to 0 ⬚C, and then the temperature stalls for nearly an hour because the heat absorbed from the air goes toward melting the ice rather than to heating the sample. When the ice completely melts, the temperature again starts to rise from 0 ⬚C to room temperature. For the propane clathrate sample the temperature does not stall at 0 ⬚C, indicating that the sample contained little pure ice. Instead, the clearest indicator of dissociation is the pressure release. As the pressure increases, the dissociation temperature also increases and the temperature holds between 0 ⬚C and 7 ⬚C for almost two hours. Plotting the pressure versus the temperature for the clathrate hydrate yields a phase line for comparison to the phase diagram. Enthalpy of Dissociation Once the pressure and temperature data of the propane hydrate have been recorded, these can be used to calculate the molar enthalpy of dissociation, ∆dH, by employing the Clausius–Clapeyron equation. The procedure is described in detail in the Supplemental Material.W Melting Point Elevation and Depression Most introductory chemistry textbooks discuss melting point depression in some detail. Here we have an example where addition of a nonpolar gas to water raises the apparent melting point. Students can be asked to explore this phenomenon by measuring the melting points for several mixtures. Even for simple salts, such as NaCl and CaCl2, melting point depression experiments are not trivial to perform and understand. We offer suggestions for further laboratory work to explore these issues. Also, we describe a procedure for making a THF hydrate clathrate that does not require the use of pressure cells and a rudimentary synthesis of the original chlorine hydrate using bleach and an HCl solution. Procedures for each of these experiments, including safety precautions, are described in the Supplemental Material.W
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Figure 2. Temperature versus time for a melting snow sample and temperature and pressure versus time for the dissociating propane clathrate-hydrate. See text for details.
An experiment, suitable for high school and undergraduate students, for preparing and studying propane hydrate has been developed and is presented here. These experiments introduce students to a surprising but important water-based, hydrogen-bonded solid. This intriguing solid, formed from two chemicals that are immiscible in the liquid state, opens many opportunities for discussing the interplay of intermolecular forces, thermodynamics, and solid structure. Also, it informs students about an immense natural reservoir of methane. Acknowledgments This work was supported by a National Science Foundation Collaborative Research in Chemistry Grant, Award No. 0404743. We also thank Joanne Abbondondola and Minhtam Vu for help in the experimental design and Laura Stern for explaining her clathrate synthesis techniques. The sample cells were designed and fabricated by Lee Moritz. One of the reviewers made numerous suggestions that resulted in significant improvements to this manuscript and the Supplemental Material.W WSupplemental
Material
Detailed background information, student procedures, sample questions, and instructor’s notes, including drawings for fabricating the sample cells and a discussion of two component phase diagrams, are available in this issue of JCE Online. Literature Cited 1. Sloan, E. D. Clathrate Hydrates of Natural Gases, 2nd ed.; Marcel Dekker: New York, 1998. 2. Davy, H. Philos. Trans. R. Soc. London 1811, 101, 1–35. 3. Sloan, E. D. Ind. Eng. Chem. Res. 2000, 39, 3123–3129. 4. Mao, W. L.; Mao, H. Proceedings of the National Academy of Sciences of the United States of America 2004, 101, 708–710. 5. Koh, C. A. Chemical Society Reviews 2002, 31, 157–167. 6. Handa, Y. P.; Tse, J. S. J. Phys. Chem. 1986, 90, 5517–5921. 7. Stern, L. A.; Kirby, S. H.; Durham, W. B. Science 1996, 273, 1843–1848.
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