Welcome to the web page for Ken Janda s laboratory! Currently, my students and I are studying a class of solids called gas clathrate hydrates. In these species, a water lattice is formed that creates cages, and these cages then enclose guest species that make the whole structure stable under certain conditions. Formally, these are called solid solutions since there is no formal chemical bonding between the guest molecules and the water lattice, and they are only stable as solid crystals. You may have heard about methane hydrates that are found in ocean and deep lake sediments and the arctic permafrost. They are one very important example of a gas clathrate hydrate. In the video above, we demonstrate the combustion of a propane hydrate clathrate. The sample looks like a cylinder of packed snow. However, when ignited, the ice lattice melts and the caged propane burns. The propane density in the sample is similar to that in a gas cylinder at 180 atmospheres of pressure. Gas clathrate hydrates are a fascinating for both theoretical and practical reasons. The methane hydrates mentioned above could potentially be an important source of energy for our modern economy. Burning methane contributes to global warming, but much less than does burning oil or coal. By some estimates, there is enough methane stored in ocean sediments to power our economy for 500 years! Gas hydrates could also provide a safe method to ship methane and LPG around the world economically and safely. However, methane hydrate is also a potential source of disaster. Methane is a powerful greenhouse gas, more dangerous than carbon monoxide. If the natural methane hydrates are released as global warming progresses, the heating could accelerate out of control with drastic results. The Deepwater Horizon disaster in the Gulf of Mexico in 2010 may well have been triggered by drilling through a methane hydrate deposit. Explore the Gas Hydrates link on this page to learn more about these issues.
The theoretical interest in gas hydrate clathrates stems from the fact that they are solids at near ambient temperatures and pressures, yet their stability is due to weak van der Waals forces. These species were first observed by Humphrey Davy in 1811. He was trying to produce solid chlorine. However, since he was working with an impure sample, he made a green solid that was stable to 283 K. This is now known to be chlorine clathrate hydrate. Although water has thirteen different solid phases, none of them resemble those of they hydrate lattices. The weak van der Waals interactions between the guest molecules and the water lattice are just enough to make the solid solution stable. This makes computer simulation of these species difficult since the computer must correctly calculate the chemical bonds of the water and guest molecules, the hydrogen bonds between the water molecules, and the van der Waals forces between the guest and water molecules.
Currently we have two active research projects. Several of my students are studying the kinetics of gas hydrate formation from ice particles and the potential guest gas. We have learned, for instance, that propane reacts much more quickly with small ice particles than with large ones. This is not surprising. The big surprise of this work is that the reaction speeds up as the temperature is lowered below 273 K. Most chemical reactions slow down as the temperature goes down. This one speeds up! (For more information, see: Journal of Physical Chemistry C, Vol 116, page 19062.) I am especially proud of the paper because all of the work, and most of the writing, was performed by a single U.C. Irvine undergraduate student, Joel Rivera. I mainly work with undergraduate students now that I am a Dean. For more on previous gas clathrate hydrate work that my students performed, please explore the links on this page. The second current project is the one being performed in collaboration with Professor Martin s group. We are using solid- state NMR techniques to study the motions of water molecules in the hydrate lattice and guest molecules within the lattice. This work is still in the preliminary stages, and we need to get more results before publically discussing them. A little background: I have always tried to choose research projects that are just beyond the ability of current chemical theory. The idea is to measure something that gives theoreticians a goal, and a test for the theories they are developing. For forty years I have been fascinated by the forces that hold molecules together and make them interact with each other. During this time, the ability of scientists to measure and compute the structure and dynamics of molecules has improved dramatically. During my Ph.D. studies, I was the first person to learn how one hydrogen chloride molecule and one hydrogen fluoride molecule would stick together. (For more information, see: Journal of Chemical Physics, 67, 5162 (1977) At the time, this
phenomenon was too difficult for chemical theory to predict. Ten years later, we were able to stick a single helium atom onto a single chlorine molecule and excite the van der Waals molecule to a single quantum state with a laser. We were then able to measure exactly how the laser energy caused the molecule to dissociate: how long the process took and how the excess energy was distributed among the product translational, rotational and vibrational motions. We then worked with an international group of scientists to reproduce the experimental results on a computer. (For more information, see: Journal of Chemical Physics, 89, 3535 (1988) Today, either of these projects would be so easy that an experiment would be almost unnecessary. The results can be calculated more easily than measured. However, only slightly more complicated problems are still impossible to predict by theory. Only in 2011 were we able to measure in detail the dissociation dynamics for a van der Waals cluster composed of two neon atoms and a bromine molecule. (For more information, see: Journal of Chemical Physics, 132 (22), 221103 (2010). There are several other previous research themes of which I am still very proud. As a postdoctoral fellow I collaborated on a project in which our team was able to bounce argon and xenon atoms off of single crystal surfaces. We were able to measure the probability that the atoms would stick to or bounce off the surface, and many other subtle details of the collisions. (For more information, see: Physical Review Letters, 43, 1175-1177 (1979), Journal of Chemical Physics, 78, 1559
(1983), and Journal of Chemical Physics, 83, 1376 (1985). Later, my collaborators and I were able to measure the kinetics of recombinative desorption of hydrogen molecules from a silicon surface covered with hydrogen atoms. (For more information, see: Physical Review Letters, 62, 567, (1989) and Journal of Chemical Physics, 92, 5700 (1990.) Another fascinating topic for me was the behavior of liquid like He clusters. Helium is the only substance that never freezes at 1 atm, even at absolute zero temperature. So, understanding how several to several thousand helium atoms behave at low temperatures is lots of fun. One of the most technically demanding projects I ve been associated with is understanding in great detail the quantum states and dissociation dynamics of a van der Waals cluster made up of two helium atoms and a chlorine molecule. (For more information, see: Journal of Chemical Physics, 95, 729 (1991) and Journal of Chemical Physics, 113, 7252-7267 (2000).) My students and collaborators also performed a series of fascinating studies that showed what happens after an atom is ionized inside of a large helium cluster. (For more information, see: Journal of Chemical Physics, 108, 9371-9382 (1998), Journal of Chemical Physics, 108, 9351-9361 (1998), Journal of Chemical Physics, 109, 10195-10200 (1998), Journal of Chemical Physics, 109, 10679-10687 (1998) and Journal of Chemical Physics, 109, 10873-10884 (1998). Finally, I would like to mention an experimental and theoretical tour- de- force that is a mostly unrecognized treasure. My students and collaborators performed spectroscopy on the iodine chloride molecule with 1 part per 500,000,000 resolution. They were able to measure the effects due to the fact that neither the iodine atom, nor the chlorine atom nuclei are spherical. This asymmetry allowed us to measure the asymmetry of the molecules electronic wave function as a function of the distance between the two nuclei. In essence, we measured the rehybridization of the electronic wave function as the molecule vibrates. I think it is fair to claim that this is the most detailed study of a heavy molecules wave function currently in the literature. (For more information, see: Journal of Chemical Physics, 101, 7221 (1994), Journal of Chemical Physics, 103, 9125 (1995) and Journal of Chemical Physics, 113, 7211-7223 (2000). To date, ab initio electronic structure theory has yet to be able to match these experimental results.