Pulling nitrogen from the air to make fertilizer is one of the biggest commercial chemical enterprises in the world, involving a complex process of heating dangerously combustible hydrogen under very high pressures. Meanwhile, in the ground just outside these nitrogen-fixation factories, tiny bacteria are also pulling nitrogen from the air-but they are doing it at room temperature and at everyday pressures while being every bit as efficient as the factories. Now a team of researchers thinks it's found a key piece of the puzzle that lets these microorganisms best the world's biggest chemical production business.
"Nature figured out how to fix nitrogen a billion years before we did," says Patrick Holland, assistant professor of chemistry at the University of Rochester and author of the research published in the latest issue of the Journal of the American Chemical Society. "We're just playing catch-up." Bacteria called azotrophs on the roots of plants take nitrogen from the air and "fix" it, turning it and hydrogen into ammonia that plants use to make DNA and proteins. Animals get nitrogen for their DNA by eating plants, so the very basis of most life on Earth depends on a few bacteria living on the roots of plants.
Since the early part of the 20th century when chemist Fritz Haber discovered that iron can be used on a large scale to fix nitrogen, iron has been recognized as playing an important role as a catalyst. Bacteria also use iron, but Holland has found that the way in which bacteria's iron is bonded may be the key to how nature can fix nitrogen without the pressures and temperatures the man-made process demands.
Holland originally set out to investigate the fundamental chemistry of iron compounds that have three atoms attached to each iron, but an unexpected chemical reaction shed light on nature's secret nitrogen formula. Since the compounds he was studying would likely react with air and water, Jeremy Smith, a postdoctoral fellow working with Holland, worked with the compounds in sealed boxes full of a gas that is usually considered safe because it rarely reacts with anything-nitrogen.
Not yet realizing that the nitrogen had reacted with the iron, Holland's team used X-ray diffraction, in which a computer-controlled device bombards the compound with X-rays, to inspect its structure. The X-rays diffracted around the atoms like sunlight bending through the crystals in a chandelier. A computer scans all the diffractions to create an image of the atoms' layout. The complex process is like trying to map the position of crystals in the chandelier by looking only at their scattered reflections.
When the computer had finished mapping the results, it had a surprise waiting.
"When Jeremy and I looked at the analysis, we saw a nitrogen molecule stuck to iron and being stretched apart," says Holland. Since the iron atoms in his model were bonded in only three places instead of the more usual six, "three-coordinate iron seems to be especially good for holding nitrogen tightly."
The iron atoms bind to a molecule of N2 (two atoms of nitrogen) floating in the air and stretch it, weakening the bond between the two nitrogen atoms. This makes the nitrogen more likely to link up with hydrogen to create ammonia (NH3)-the useful form with one nitrogen atom combined with three hydrogen atoms.
"We've shown that this type of iron can stretch nitrogen bonds, but now we want to see how this enzyme goes to the next step and breaks that bond," says Holland. Knowing how Nature so easily pulls nitrogen from the atmosphere may lead to a better understanding of how the ecology may balance itself, and may eventually provide a safer, more economical method of nitrogen fixing for the multi-billion-dollar fertilizer industry.
The research was funded by the University of Rochester.
The above post is reprinted from materials provided by University Of Rochester. Note: Materials may be edited for content and length.
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