At the bottom of the world’s oceans, ice studs the rock just below the water. But the ice isn’t made just from water; the ice-like compounds are methane clathrates, cage-like structures of water molecules that form around a guest molecule. As the quest for alternative energy sources continues, scientists and engineers are exploring new options, and methane clathrates are a distinct possibility. But, if released, they could potentially be harmful to the atmosphere; whether these molecules are good or evil has yet to be seen.
Discovered only recently, clathrates are not yet fully understood by scientists. Clathrates are ice-like structures around a guest molecule, which can be gas, liquid or solid, but the cages do not actually bond with the guest. This means that the guest molecule is trapped in the cage, but it moves about uninhibited once the cage is compromised. In an effort to better understand the clathrates’ behaviors, Kristen Pallen ’12, Lennox Chitsike ’13, Aaron Danilack ’13 and Pauline Wafula ’13 will be testing the effects of salts on these mysterious clathrates with the guidance of Assistant Professor of Chemistry Camille Jones.
The team will be placing multiple different guest molecules inside the clathrate structure; Pallen and Danilack are currently testing tetrahydrofuran, Chitsike is working with Trimethylene oxide, and Wafula is testing cyclobutanone and propylene oxide. Different guest molecules cause the clathrate structure to form around them at different temperatures and pressures.
But the team will also use different salts to see how they affect the stability of the clathrates; although clathrates are relatively new in the world of chemistry, an even more groundbreaking concept is that salt could actually stabilize the ice instead of causing it to melt. While Pallen and Chitsike are working with fluoride salts, Danilack has been testing the effects of chloride salts (such as sodium chloride, potassium chloride, caesium chloride, and lithium chloride) on the clathrates’ melting points.
“Salts are generally considered clathrate inhibitors,” Pallen said. “In theory, they should make them less stable. But we’ve noticed that [the salts] are making them somewhat more stable, particularly fluoride salts.”
Think of an icy road during the winter. It is common practice to put salt on the road because the salt lowers the ice’s melting point, keeping the ice in its liquid state even if the temperature is below freezing. Scientists generally believe that salts would make ice structures less stable, like with the ice on the roads. But some salts have actually been observed to make the cage structure of a clathrate stronger by forming bonds with other molecules or even substituting or replacing another molecule in the lattice.
Chitsike is working to correlate the melting point with the pH of his Trimethylene oxide hydrate when combined with a sodium fluoride salt. So far, he has found that the salt is slightly alkaline because, when the components of the salt dissociate (or break down), they form hydroxide ions in water. “This finding is crucial because it’s going to enable me to evaluate the degree of [the salt’s] dissociation in solution and ascertain the concentration of fluoride ions present,” Chitsike said.
But still, he has found that these dissociated components of salts can raise the clathrate’s melting point, making it more stable: “If that’s true that some salts like sodium flouride stabilize hydrates, then we might conclude that some hydrates are actually reservoirs for anions and cations of the salts that stabilize them,” Chitsike explained. This means that, after the ions dissociate, they find their way into the clathrate structure, either taking the place of or bonding to other ions in the compound.
But Danilack’s chloride salts act differently. “Chloride ions are too large and don’t hydrogen bond well, so unlike the fluoride ions, the chloride cannot become part of the hydrate structure,” he said. So far, Danilack has found that sodium chloride always lowers the clathrates’ melting point more than does potassium chloride when comparing equal concentrations of both salts in a solution—an unexpected conclusion, as he hypothesized that different chloride salts would only change the melting point depending on their concentration, not the other atom bonded to the chloride ion.
By better understanding the clathrates’ behaviors as well as the effect of salts, scientists can better harness that energy into a potentially feasible energy source, or better understand it if it proves more harmful than beneficial.
The group's summer research was funded through the Edward and Virginia Taylor Fund for Student/Faculty Research in Chemistry, established in 2008 through a gift from Ted ’46 and Virginia to inspire students interested in chemical research and to facilitate their work with outstanding faculty.
Discovered only recently, clathrates are not yet fully understood by scientists. Clathrates are ice-like structures around a guest molecule, which can be gas, liquid or solid, but the cages do not actually bond with the guest. This means that the guest molecule is trapped in the cage, but it moves about uninhibited once the cage is compromised. In an effort to better understand the clathrates’ behaviors, Kristen Pallen ’12, Lennox Chitsike ’13, Aaron Danilack ’13 and Pauline Wafula ’13 will be testing the effects of salts on these mysterious clathrates with the guidance of Assistant Professor of Chemistry Camille Jones.
The team will be placing multiple different guest molecules inside the clathrate structure; Pallen and Danilack are currently testing tetrahydrofuran, Chitsike is working with Trimethylene oxide, and Wafula is testing cyclobutanone and propylene oxide. Different guest molecules cause the clathrate structure to form around them at different temperatures and pressures.
But the team will also use different salts to see how they affect the stability of the clathrates; although clathrates are relatively new in the world of chemistry, an even more groundbreaking concept is that salt could actually stabilize the ice instead of causing it to melt. While Pallen and Chitsike are working with fluoride salts, Danilack has been testing the effects of chloride salts (such as sodium chloride, potassium chloride, caesium chloride, and lithium chloride) on the clathrates’ melting points.
“Salts are generally considered clathrate inhibitors,” Pallen said. “In theory, they should make them less stable. But we’ve noticed that [the salts] are making them somewhat more stable, particularly fluoride salts.”
Think of an icy road during the winter. It is common practice to put salt on the road because the salt lowers the ice’s melting point, keeping the ice in its liquid state even if the temperature is below freezing. Scientists generally believe that salts would make ice structures less stable, like with the ice on the roads. But some salts have actually been observed to make the cage structure of a clathrate stronger by forming bonds with other molecules or even substituting or replacing another molecule in the lattice.
Chitsike is working to correlate the melting point with the pH of his Trimethylene oxide hydrate when combined with a sodium fluoride salt. So far, he has found that the salt is slightly alkaline because, when the components of the salt dissociate (or break down), they form hydroxide ions in water. “This finding is crucial because it’s going to enable me to evaluate the degree of [the salt’s] dissociation in solution and ascertain the concentration of fluoride ions present,” Chitsike said.
But still, he has found that these dissociated components of salts can raise the clathrate’s melting point, making it more stable: “If that’s true that some salts like sodium flouride stabilize hydrates, then we might conclude that some hydrates are actually reservoirs for anions and cations of the salts that stabilize them,” Chitsike explained. This means that, after the ions dissociate, they find their way into the clathrate structure, either taking the place of or bonding to other ions in the compound.
But Danilack’s chloride salts act differently. “Chloride ions are too large and don’t hydrogen bond well, so unlike the fluoride ions, the chloride cannot become part of the hydrate structure,” he said. So far, Danilack has found that sodium chloride always lowers the clathrates’ melting point more than does potassium chloride when comparing equal concentrations of both salts in a solution—an unexpected conclusion, as he hypothesized that different chloride salts would only change the melting point depending on their concentration, not the other atom bonded to the chloride ion.
By better understanding the clathrates’ behaviors as well as the effect of salts, scientists can better harness that energy into a potentially feasible energy source, or better understand it if it proves more harmful than beneficial.
The group's summer research was funded through the Edward and Virginia Taylor Fund for Student/Faculty Research in Chemistry, established in 2008 through a gift from Ted ’46 and Virginia to inspire students interested in chemical research and to facilitate their work with outstanding faculty.
Aaron Danilack is a graduate of Crestwood High School, Mountain Top, Pa.; Pauline Wafula graduated from Loreto High School in Limuru, Kenya; Kristen Pallen is a graduate of Trinity College School, Port Hope, Ontario, Canada; and Lennox Chitsike graduated from St. Faith's High in Rusape, Zimbabwe.
