Getting A Cheaper Nitrogen Fix  

Professor Richard Schrock from MIT has an article on research into replacing the Haber Bosch process for producing ammonia (and thus fertiliser).

Molecular nitrogen (dinitrogen, N=N) makes up about 78 percent of the atmosphere. It is the most unreactive diatomic species known. Interestingly, however, nitrogen is required for all life; it is used to build proteins and DNA. Therefore, dinitrogen must be turned into a molecule that can be assimilated readily by plants. That molecule is ammonia, NH3.

Prior to World War I, the iron-catalyzed Haber-Bosch process for ammonia synthesis at high temperatures (350 to 550 °C) and pressures (150 to 350 atmospheres) from dinitrogen and dihydrogen (H2) was discovered. It is perhaps the most important industrial process ever developed and responsible for a dramatic increase in the population of the earth during the 20th century, because it supplies a reliable source of nitrogen for fertilizers. But because the Haber-Bosch process requires high temperatures and pressures, it consumes tremendous amounts of energy; it is estimated that as much as 1 percent of the world's total energy consumption is devoted to the process.

Nature also reduces dinitrogen using metalloenzymes in bacteria and blue-green algae, but at only one atmosphere of pressure and mild temperatures. The metalloenzymes, called nitrogenases, contain iron and usually molybdenum. Ever since their discovery more than 40 years ago, chemists have speculated about how reduction of dinitrogen occurs and whether an "artificial" nitrogenase could be developed that would lead to a more energy-efficient process than Haber-Bosch. Perhaps a thousand man-years and billions of dollars have been spent studying how nitrogenases work and trying to make artificial ones.

In 2003, my group showed that it is possible to make ammonia catalytically from dinitrogen, protons, and electrons. This is accomplished at a single molybdenum metal center. In the presence of protons and electrons in a non-aqueous medium, dinitrogen is reduced to ammonia with an efficiency in electrons of around 65 percent; the remaining electrons are used to make dihydrogen, which is in this context a wasteful and undesirable product. Our catalyst is not great, but it is a start.

Nature has developed a highly optimized version of the nitrogen reduction process over a period of a few billion years. Ours is an "artificial" nitrogenase that is barely catalytic. We are trying to identify the key problem or problems that prevent it from working well. Perhaps then we can improve its efficiency.

Can we design catalysts that will be as efficient as natural nitrogenases? Possibly. Will the Haber-Bosch process ever be replaced by catalysts that do not operate at high pressures and temperatures? Unknown. Only time, money, and ingenuity will reveal the answer.

Combine air and natural gas over an iron oxide catalyst under high pressure and intense heat and what do you get? The answer, surprisingly, is plant food: ammonia, the chemical precursor to nitrogen fertilizers.

Ammonia gets converted into nitrites and nitrates, which when sprinkled onto plants, allow them to grow larger. This is the basic idea behind the huge increases in agricultural yields, doubling between 1950 and 1990, seen in the 20th century. (Caveats about the "quality" of this growth and the environmental impacts of nitrogen are noted, but left aside for a later post in this continuing series).

Back around 1915, the world produced almost no nitrogen fertilizer, largely because there was no usable nitrogen supply. Now, the world produces about 87 million tons of N-based fertilizers. This increase is primarily due to the Haber-Bosch process for pulling nitrogen out of the air. (The development of new plant varieties that are able to soak in excess nitrogen will also be the subject of a separate post).

Clearly, the Haber-Bosch process has been successful. As we've noted before, at least one professor has estimated that 40 percent of the world's food can be traced back to the process. But the process is encountering major problems in the increasingly resource-constrained world.

Here's why: the main reaction in the process is cooking N2 and H2 together at 500 degrees Celsius and 200 atmospheres of pressure. You need all that heat and pressure because breaking apart an N2 molecule turns out to be incredibly difficult. A nitrogen atom has five electrons in its outer shell (valence electrons), so it has a tendency to share three electrons with another nitrogen atom to reach its stable (octet rule) state. That's what generates dinitrogen's triple covalent bond, one of the strongest in nature. The energy required to break the bond is 946 kilojoules of energy per mole of nitrogen, or twice the energy required to bust an O2 molecule.

Luckily, or so we thought, fossil fuels were cheap, widely available, and incredibly energy dense: 1 cubic foot of natural gas contains 1.055 gigajoules of energy. That's enough energy to convert a lot of moles of nitrogen into ammonia. So, once the Haber-Bosch process established it could be done, chemists across the world began to burn a lot of natural gas to get dinitrogen to react with hydrogen. And where do we get the hydrogen? Why, we use the natural gas for that too, naturally: it is CH4 after all.

Taken together, there's a lot of natural gas going into the production of nitrogen fertilizer. So much so that when I tweeted about my fertilizer investigation, my friend Celeste LeCompte, managing editor at the Sustainable Industries Journal, tweeted back, "Think: natural gas."

In effect, we've been pumping fossil energy into our food supply, and eating it. While diminishing fossil fuel supplies and climate concerns have given us perfect hindsight into why this could be a dubious path for the future, at the time, it must have seemed like an excellent idea, given that the alternative--not producing enough food--was both real and horrific.

Until relatively recently, the price of natural gas, which tracks the price of oil very closely, was relatively low. Now, with oil over $120 a barrel and natural gas prices having doubled since the mid-90s to over $11 per thousand cubic feet of the stuff, the cost of ammonia has tripled. As in biofuels or alternative energy, the rising cost of oil is driving innovation.

As we've noted before, legumes developed symbiotic relationships with bacteria who can pull nitrogen out of the air at room temperature and standard atmospheric pressure. They use a specialized enzyme known as a nitrogenase that consists of iron and the metal molybdenum. In fact, scientists estimate that 200 million tons of nitrogen are fixed via natural processes, or more than twice human production.

Now teams of scientists across the world from Richard Schrock at MIT to David Tyler at the University of Oregon are racing to find just the right catalyst to recreate the natural nitrogen fixation process. While they wouldn't eliminate the use of natural gas as a feedstock, they would reduce the amount of energy used in the creation of ammonia. How much? Eliminating the Haber-Bosch process, which uses an estimated one percent of the world's total 15 terawatts of energy consumption (xls) would mean 150 gigawatts of energy savings for the world. That's about as much coal generating capacity as the US is planning to add between now and 2030.