6
Life, Death, and Bread from the Air
It is through the laboratory that starvation may ultimately be turned into plenty … The fixation of atmospheric nitrogen is one of the great discoveries, awaiting the genius of chemists.
—Sir William Crookes
There is no evil in the atom; only in men’s souls.
—Adlai Stevenson
Atoms are what bombs and toxins are made of, and many people fear them because of it. But the most wonderful things in life are also made of atoms, as are you and I. Bees can make delicious honey although they can also sting, and people may seem to be satanic or saintly when in fact both traits blend within all of us. This multifaceted nature of things is nowhere more vividly personified than in the life of a German chemist named Fritz Haber.
The inventor of a way to create explosives and fertilizer from nitrogen in the air, Haber contributed to the deaths of millions, possibly including the suicides of his wife and son, but he also contributed to the sustenance of billions over the course of the last century. For his invention and its industrial applications he shared a Nobel Prize in 1918, an award whose endowment, ironically, was based on a fortune made by Alfred Nobel through the production of nitrogen-based explosives.
The legacy of Fritz Haber revolves around nitrogen, a vital atomic element that shares his multifaceted nature. It comprises more than three-quarters of every breath you draw, every sound you hear, and every word you speak. Sometimes it helps to acidify lakes or reduces buildings to rubble. Through the ages it has linked people, wildlife, and plants to dirt-dwelling bacteria, thunderbolts, and the iconic blue of the sky. And, perhaps most astoundingly, it now links much of the atomic framework of your body to Fritz Haber.
* * *
Look at the sky on a clear sunny day, as you have surely done many times before.
Now really look at it.
How strange it is. At first glance, you might think that the sky is a solid ceiling as the ancients did. When the book of Genesis describes the “firmament” of the heavens, it really means “firm.” But of course there is nothing so solid about that blue sea of air. What you’re seeing is a luminous haze, a diffuse and permeable mist that encases the illuminated half of the earth like a fogged contact lens. The sun shines through it, meteors penetrate it without leaving holes, and when a bird, cloud, or jet passes you see it clearly in the transparent medium overhead. So where does the blueness come from?
This question probably dates back to the earliest days of language, but the modern scientific answer to it crystallized only recently around the work of many investigators including another Nobel laureate, Albert Einstein. Known as both a peace activist and as the father of the atomic bomb, Einstein made numerous discoveries that are less widely known than his research on relativity theory. In 1910, while his friend Haber was working on nitrogen as a source of ammonia, Einstein published a paper that helped to reveal nitrogen’s role in creating the color of the sky.
To understand that role, it is important to remember that nitrogen comprises 78 percent of the gas in a typical parcel of air. As with the oxygen gas that makes up most of the remainder, these molecules are paired atoms of the same name, but the twin atoms in a nitrogen gas molecule—or dinitrogen—look and act differently from the other components of the atmosphere.
When you draw your next breath, some dinitrogen will diffuse into your body along with the oxygen. If you are a scuba diver, you know that breathing pressurized air under water puts extra nitrogen gas into your bloodstream, which can bubble out of solution if you surface too rapidly without decompressing. Such bubbles can cause severe joint pain, block circulation, and damage nerves in a potentially fatal condition known as “the bends.” But under most circumstances your cells ignore nitrogen molecules as irrelevant transients. Although you need nitrogen atoms to build and maintain your body, you can’t simply inhale your nitrogen supply as you do your oxygen. You have to eat it, and to do that you need the help of other living things in order to make it edible.
The problem is that nitrogen gas is to biological nitrogen as the dilute nectar of a flower is to honey. It requires processing to become most useful to you. You could die from a shortage of nitrogen atoms in your diet while carrying a full load of nitrogen gas in your lungs. Powerful tools are needed to convert that gas into more accessible forms, and only a handful of species possess them. These lucky few hold the keys to an inexhaustible trove of atoms that you and every other creature on Earth can’t live without.
The feature that makes this element so hard to get also contributes to the color of the sky. Nitrogen seems to prefer the company of its own kind, at least when you come across it in its most common form. Other atmospheric gas atoms can also pair up in similar fashion by sharing one or more of their electrons in covalent bonds, but not nearly so tightly. The twin atoms in a hydrogen molecule are joined by one bond, and those in an oxygen molecule by two. But in dinitrogen the paired atoms clamp together with such a tenacious threefold grip that it becomes extremely difficult to separate them once they get hold of one another. If close friends can be said to be “joined at the hip,” then the atoms in a nitrogen molecule are more like passionate lovers who are locked together from head to toe.
The electron clouds that surround and bind the atoms in air molecules such as dinitrogen are not perfectly rigid. If you could give such a molecule a vigorous shake the shell of electrons surrounding it would jostle slightly, and this responsiveness makes a remarkable thing happen when the sun rises on a clear morning. As the bees in the meadows of the world awaken and begin the day’s harvest, the air molecules above them begin to hum as well, though not in waves of sound. Instead, the air buzzes with light.
The daytime color of Earth’s nitrogenous sky is more than a mere backdrop to scenery: It also affects you physically. Recent investigations into the effects of light on human physiology show that exposure of your retinas to blue light can suppress your production of melatonin, a hormone that induces sleep. In days past this response to sky-filtered sunlight probably helped to keep your early ancestors alert and active between dawn and dusk through an unconscious, reflexive sensitivity that sleep researchers call “sightless vision.” Today, however, some experts believe that the blue-enriched glows of computer screens, televisions, and other forms of artificial illumination that now shine around the clock may contribute to a wide spectrum of ailments from depression and insomnia to fatigue-related accidents.
To illustrate how the blueness of a sunlit sky is made, imagine the molecules in the air around you as a swarm of invisibly small honeybees—friendly honeybees, of course. Among them you’ll find three kinds of movement going on at the same time. Winds blow the whole swarm here and there, and the individual bees also dart back and forth in a chaotic thermal dance. But much of the color of the sky emerges from the air molecules themselves, as the humming of bees emerges from their bodies.
The buzz of a bee comes from the rapid-fire shivering of tiny muscles which can power the beating of wings, shake pollen loose from a blossom, or warm the insect on a chilly morning. If you allow a good-natured bee to rest on your hand you can often feel the tingle of those vibrations. The atmosphere radiates light as bees radiate sound, but it does so with the help of electrons rather than muscles. When sunlight strikes dinitrogen, the molecule’s electron cloud twitches rhythmically in response, and physicists who study that silent optical buzz sometimes refer to air molecules in essentially musical terms as “harmonic oscillators.” One expert recently described the phenomenon to me, along with a strict warning not to bungle the explanation.
“Most people get this wrong,” Craig Bohren told me, “even many scientists. As long as you don’t make the mistake of saying that the sky is only blue or even predominantly blue, you are on the side of the angels.”
With that surprising admonition, this well respected atmospheric scientist, now retired from Pennsylvania State University, led me through the basics of sky color.
“The nitrogen and oxygen molecules that dominate the atmosphere are much smaller than light waves, and this allows them to scatter sunlight in every direction when it hits them. They actually scatter all kinds of colors, but most notably the ones with shorter wavelengths.” That would be blue, but also violet. So—why isn’t the sky closer to violet?
“Well,” he said with a chuckle, “in a sense it actually is, but you don’t see it that way. Your eyes aren’t as sensitive to violet as they are to blue, and your brain doesn’t analyze light like a spectrophotometer would. What you see above you is partly of your own making, and you just give the name ‘blue’ to what you perceive up there.”
According to Bohren’s measurements, only about one-fifth of the light in a bright noontime sky is actually blue. “Even Einstein overlooked this detail when he investigated optical scattering, because he never looked at a sky-light spectrum.”
And how, exactly, does an air molecule scatter light?
“If you’re sure you want to get into it,” he began after a pause, “it has to do with how the electron cloud responds when a rapidly oscillating light wave strikes it. The motion of the electrons relative to the nucleus generates an electromagnetic disturbance whose wavelength or color is related to the wavelength of the incoming light. An air molecule does this with many light waves at once, but its scattering of the blue and violet light is greater than that of, say, red or yellow.”
In other words, molecular nitrogen can scatter the full palette of colors that stream out of the sun, much as a cell phone can reproduce many of the sounds of a symphony that a friend might send to you surreptitiously from a concert hall. But like the small speaker in your phone, which emits the higher pitches of a violin section more effectively than the deep booming of a bass drum, the tiny molecules of the atmosphere broadcast the relatively shorter waves of blue and violet light with greater fidelity. We are a bit tone-deaf when it comes to the high violets, but we can still get plenty of pleasure out of those equally glorious blues.
Multiply this process trillions of times per second throughout the lower atmosphere, add the limitations of human vision, and you end up perceiving something akin to a swarm of tiny, luminescent blue bees. The effect is so powerful that it can block your view of the stars during daytime, and you can read by scattered sky-illumination alone for quite some time after sundown. What appears to be an opaque dome overhead is more than just a passive ceiling: it glows with its own atomic light. But in the end, it is up to you to make it blue.
* * *
Although nitrogen atoms comprise only 3 percent of your total body mass, they are key components of molecules that make you look and act as you do. The carbohydrates and lipids in your body consist mostly of the three elements that comprise carbon dioxide and water, but in order to make the thousands of kinds of protein that keep you alive, you must also add nitrogen to the mix. Nitrogen atoms represent 10 to 15 percent of the dry mass of your muscles, and four nitrogen atoms cradle each rust-colored iron atom in the hemoglobin of your blood. All your enzymes, antibodies, and genes contain them, as do the ion pumps on your neurons and the cartilage in your nose. Without nitrogen in your diet, the most substantial parts of you that could be built from carbon, hydrogen, and oxygen alone would be your body fats and fluids.
The electron clouds that make nitrogen gas respond to sunlight also require a cell to have the molecular equivalent of a wrench in order to pry it apart and make other things from the atomic pieces. Luckily for us, some organisms do possess such tools. Microbes embedded in the roots of alder shrubs perform some of that work, but chief among the world’s nitrogen fixers are colorful cyanobacteria that live in plankton or embedded in the crinkly tissues of lichens, and soil bacteria that colonize the roots of legumes such as alfalfa, clover, and beans. Because of these alliances among bacteria and their hosts, planting a field with alfalfa is like spraying it with fertilizer.
The molecular wrenches of nitrogen-fixing bacteria are iron-bearing enzymes called nitrogenases. A nitrogenase can snap dinitrogen in half, then decorate each atom with three hydrogens, thereby forming biologically useful ammonia. Unlike typical chemicals that are consumed or neutralized in reactions, enzymes such as nitrogenases are sturdy implements that perform their tasks over and over for as long as energy and raw materials are available.
Two potential sources of your nitrogen atoms. Left: Mimosa pudica, a tropical legume that houses nitrogen-fixing bacteria in its root nodules. Right: Photomicrograph of Anabaena, a lake-dwelling cyanobacterium that fixes nitrogen in enlarged cells such as the one in the center of this strand. Photos by Curt Stager
Through the ages nitrogen-fixing bacteria such as these have monopolized the production of nitrogen compounds that the rest of life on Earth has craved, stolen, and killed for. If they had been human traders of such a commodity, they would have been a fabulously wealthy cartel. Instead, root nodule bacteria have limited their demands to free accommodation in cozy, subterranean lodgings, and in the case of aquatic cyanobacteria, to chemical defenses that can punish creatures that try to eat them.
Apart from bacteria, the only other noteworthy nonhuman source of useful nitrogen is lightning. A bolt of lightning may be no thicker than your thumb, but it is hotter than the surface of the sun. The intense heat tears nitrogen molecules apart and allows oxygen to cling to the separated atoms. Each air-shattering stroke leaves a vapor trail of nitrogen oxide behind, which disperses into the atmosphere and eventually falls to the ground in rain or snow where plants can get at it and feed it into various food chains. You almost certainly carry atoms from the smoke of thunderbolts inside you right now, as do most other living things.
The invisibility of most of the nitrogen that you encounter in daily life makes it difficult to develop an inituitive awareness of your connections to it. But with a little scientific information and a healthy imagination, you can still probe beneath the surface of familiar scenes and situations for deeper atomic truths. For illustrative purposes, here is one such view of the nitrogen hidden in a setting that may be familiar to you—the Kansas tornado scene in the classic film The Wizard of Oz.
As the wind picks up under a darkening sky, a ropy tentacle of cloud snakes toward a sad-looking field, and between the field and funnel stand a house and barn. As the twister approaches, the people, chickens, and horses run for cover. Press the Pause button here. What are the nitrogen atoms doing?
This portion of the movie was filmed in black and white, so you couldn’t see a blue dinitrogen sky above the long, tapered sock that makes such a convincing tornado on screen, even if it weren’t cloudy. No matter—there is plenty to consider in the motion of the air rather than its color. Most of the force of those terrible winds represents the impacts of flying nitrogen molecules against the ground and the characters as they seek shelter. The fierce flow was largely driven by the differential thermal movements of nitrogen molecules farther upwind, as the heat of the sun sped their dances in rising, swelling parcels of air. The collision of warmer, faster-dancing air masses with cooler, slower-dancing ones over Kansas spawned this devastating storm.
Flashes of lightning could be edited into those cinematic clouds as well. Every explosive bolt would leave a streak of nitrogen oxide in the turbulent air. Some of the oxides would dissolve into raindrops and soak into the soil, later to help sustain whatever crops survived the storm.
Unseen in this view are the nitrogen-rich vapors rising from the field. Bacteria in the soil have been eating leftovers from last year’s harvests and manure piles and unwinding their proteins into the gases from which they originated. Bacteria beneath an outhouse that stands behind the barn are doing the same.
Nitrogen atoms are hidden in the dry stalks in the field as well, lingering in the remains of once-green chlorophyll. They are even more abundant in the keratin feathers of the chickens as they run squawking through the barnyard.
The film resumes with Auntie Em stepping out into the gale to call for Dorothy as the dinitrogen wind whips the nitrogen-bearing keratin of her gray hair. Dorothy herself, stumbling in the blinding gusts, is too busy coordinating her nitrogen-rich muscles to hear the nitrogen-borne sound waves of her aunt’s cries. Her little dog, Toto, is faring a bit better, however, because his sharp keratin claws give him a better grip on the ground.
Similar nitrogenous connections operate all around you in the modern world, but some new ones have also entered the picture now. The spark plugs in the engines of our motor vehicles resemble miniature thunderstorms, and they fix nitrogen just as lightning does. You inhale close to three thousand gallons of nitrogen-laden air per day, on average, a volume equivalent to that of three dozen large bathtubs. A typical compact car, by comparison, inhales thirty times as much in an hour, and if you drive it 12,500 miles over the course of a year it can fix roughly eighteen pounds of nitrogen oxides. Multiply that respiratory rate by four for a large truck and by closer to a hundred for an airliner, and you can imagine why vehicle exhaust, along with the fiery innards of coal-fired power plants, has become a major component of the global nitrogen cycle.
The Canadian meteorologist Lewis Poulin estimated that traffic in and around Montreal in 2001 processed about 175 times more air per day than the city’s 1.8 million inhabitants did. But unlike the nitrogen molecules that people exhale unaltered, many of the ones that pass through those superhot engine cylinders emerge as nitrogen oxides. Scientists with the United States Environmental Protection Agency have shown that every major city on the Atlantic Coast sends a smoggy plume of nitrogen-enriched gases far out to sea on the prevailing westerly winds, along with other waste fumes. And the eastern cities, in turn, also lie downwind of similar plumes arising from cities and highways farther inland.
Roughly half of the planet’s biological nitrogen, including most of the nitrogen in your own body, is now snatched from the air by fossil fuel combustion and, above all, by a process that was first harnessed on an industrial scale to supply Germany with explosives during World War I. This artificially fixed nitrogen can now end up in anything from TNT (trinitrotoluene) to laughing gas (nitrous oxide). Even fertilizer (ammonium nitrate) has been used in the full spectrum of human endeavors from famine prevention to the Oklahoma City bombing. All these things are the legacy of one man, placing him among the most influential figures in human history.
* * *
Fritz Haber was born in 1868 in Breslau, East Prussia (now Wrocław, Poland), the son of a merchant. According to biographers, he was a brilliant scientist, a lousy husband, and an ambitious patriot. This combination of traits helps to explain much of the trajectory of his life. As he himself described his philosophy of hard work and civic duty, “I don’t want to rust out, I’d rather wear out.”
Like Einstein, Haber was a German Jew. But unlike Einstein, he converted to Christianity in order to further his advancement in an increasingly anti-Semitic society. Around 1905 he found a way to make liquid ammonia by combining nitrogen and hydrogen gas under high temperature and pressure, using an iron catalyst to push the reactions along in much the same manner that bacterial nitrogenase does. By 1909 he and the British physical chemist Robert Le Rossignol had developed a high-pressure apparatus that dripped ammonia in potentially useful quantities. Although he was already well on his way to prominence in the scientific community, perfecting this process made Haber into a public celebrity.
At that time Germany’s largest sources of nitrogen for fertilizer and explosives were layered crusts of sodium nitrate found in desert salt flats in northern Chile. Such deposits are extremely rare, and their origins remain somewhat mysterious. Sodium nitrate dissolves easily in water, so it could only accumulate to such an extent in hyperarid environments such as the Chilean desert. In a report published by the U.S. Department of the Interior in 1981, the geologist George Erickson concluded that the strange deposits had accumulated from sea spray, volcanic emissions, and the weathering of rocks and soils over millions of years. Almost any other location would have been too wet and vegetated to allow so much soluble, biologically useful nitrogen to build up in this manner. To Erickson, finding those foot-thick layers of fertilizer, some of them exceeding 50 percent purity, must have been almost as surprising as finding a huge plain of table sugar lying out in the open without it washing away or attracting hungry ants.
Whatever its origin, Chilean saltpeter (from the Latin words sal and petra, referring to salt and stone) enriched those who exploited it since the early 1800s, and by the early 1900s the United States and Britain had come to depend upon the great saltpeter flats as well. As the shadow of World War I loomed darker, however, the vulnerability of that single source and the long voyage from Chile turned Haber’s invention into a game changer for Germany.
When his fellow chemist Carl Bosch scaled the reactions up to industrial levels between 1909 and 1913, what then became known as the Haber-Bosch process turned the atmosphere itself into a gigantic nitrogen mine. Those who focused on its agricultural uses spoke in glowing terms of making “bread from the air.” But ammonia is also readily converted to nitric acid for the manufacture of explosives, and in the midst of war the new procedure couldn’t have been more desirable to the homeland. On the other hand, it also allowed the bloody conflict to continue longer than it otherwise would have if nitrogen resources had remained limited to geological deposits and root nodules.
How can a substance that promotes life when used as fertilizer also destroy it so violently? Nitrate molecules feed their own oxygen atoms to a fire so rapidly that a normally slow burn becomes a savage blast. Mix potassium nitrate with charcoal and sulfur in varying proportions, for example, and you get various forms of gunpowder that have long been used in weapons and fireworks.
When the eighteenth-century French chemist Jean-Antoine Chaptal gave nitrogen its elemental name (nitrogène, “source of nitre”), he was referring to its presence in saltpeter, which was then commonly called “nitre.” Each nitrate subunit in a molecule of potassium nitrate consists of a nitrogen atom joined to three oxygen atoms, and when a spark strikes it the nitre in gunpowder unloads its oxygen directly onto the flammable compounds around it. Then everything seems to happen at once. The flammables erupt into carbon dioxide, water vapor, and sulfurous gases, all of which expand at terrible speed along with heat and light. As the powder explosion in a gun pushes a bullet out of its barrel, nitrogen atoms that lost their oxygen baggage also fly out as nitrogen gas, where they blend back into the air from which they were originally fixed.
When you digest the proteins in your food, you eventually excrete their nitrogenous remains in the form of urea. Bacteria may later convert that urea into ammonia, which in turn, can become oxidized into nitrate suitable for explosives. During the American Civil War, the resource-starved Confederate army was sometimes reduced to making gunpowder from wood ash (for the potassium) and urea (for the nitrate). According to the Web site of the musician and author Rickey Pittman, every source of urea-based nitrate was exploited, from barns to chamber pots.
Pittman quoted a public notice posted by the government agent John Haralson as saying “The ladies of Selma are respectfully requested to preserve the chamber lye to be collected for the purpose of making nitre. A barrel will be sent around daily to collect it.” Poets quickly seized upon the tale, and one song of the day ended with this bawdy verse:
John Haralson, John Haralson, pray do invent a neater,
And somewhat less immodest way of making your saltpeter.
For ’tis an awful idea, John, gunpowdery and cranky,
That when a lady lifts her skirts she’s killing off a Yankee.
* * *
By fixing ammonia directly from the air, Fritz Haber revolutionized modern warfare and became a national hero. You might expect his celebrity status and contributions to the fatherland to have played well at home with his wife, Clara, who was herself a talented chemist. But their relationship was strained by Haber’s seeming eagerness to fulfill any demands that his country might make of him, regardless of their ethical implications.
Haber did not share Einstein’s pacifist philosophy and willingly applied his research to military ends. In his view, “A scientist belongs to his country in times of war and to all mankind in times of peace.” As the director of the Kaiser Wilhelm Institute for Physical and Electrochemistry, which is now named after him, Haber helped to develop and supervise the use of poison gas as a way to drive Allied troops from their trenches. He often said that he hoped such devastating tactics would bring a speedy victory to Germany and would therefore ultimately save lives that might otherwise be lost in a more prolonged conflict. He was unfortunately mistaken in that regard.
For Clara, her husband’s activities had begun to cross the line into complicity with evil. According to an article in Smithsonian magazine by the author Gilbert King, she protested his work in public as well as in private, calling it “barbarity” and “a perversion of the ideals of science.” He responded by accusing her of treason, further weakening their already-troubled marriage.
The last straw, it appears, came in the spring of 1915 after Haber personally supervised the first use of chlorine gas on Allied troops in Flanders. One Canadian soldier who survived later described being gassed by German chlorine as “an equivalent death to drowning only on dry land. The effects are … a knife edge of pain in the lungs and the coughing up of a greenish froth … ending finally in insensibility and death.” Shortly after hearing of the gas attack and the ghoulish suffering that resulted, Clara walked into the family courtyard, aimed a pistol at her chest, and pulled the trigger. Their thirteen-year-old son, Hermann, who found her as she lay dying, was left to grieve alone when Haber departed the next morning in order to oversee gas releases on the Eastern Front.
Left: Clara Immerwahr Haber. Right: Fritz Haber. Courtesy of Archiv der Max-Planck-Gesellschaft, Berlin-Dahlem
It is tempting to blame Haber for his wife’s suicide, though Clara left no explanation for her fatal decision. Certainly his family life left much to be desired, at least in part because of his willingness to put ambition and country above all else. But when the war ended, Haber redirected his efforts to support more peaceful goals. He spent much of the 1920s in an attempt to help Germany pay war reparations by extracting gold from seawater, and he proposed using fertilizers and pesticides to turn the Sudanese desert into an agricultural Eden.
In the end, however, the nation to which Haber had devoted himself betrayed him. After Hitler came to power in 1933, Nazi race laws began to purge Jewish scientists from even the most prestigious posts and drove Einstein and others to emigrate to America. One day the former hero of nitrogen fixation went to work at the institute only to be turned away by a porter who announced, “The Jew Haber is not allowed in here.” He resigned and exiled himself to England, where—not surprisingly—he was shunned by British scientists for his involvement in gas warfare, and then moved to Switzerland. According to the biographer Morris Goran, a friend described Haber in his postwar years as “seventy-five percent dead.” Fritz Haber succumbed to heart disease in 1934, a broken man. After his death his old friend Einstein reportedly concluded, “Haber’s life was the tragedy of the German Jew—the tragedy of unrequited love.”
Some say that it was fortunate for him that he did not live to see what arose from his work after his passing. By the cruelest of ironies, research by Haber and his associates at the Kaiser Wilhelm Institute led to the invention of Zyklon B, a nitrogen-bearing pesticide that was originally intended to protect crops but that would later be used in Nazi gas chambers to kill millions of Jews, including some of Haber’s former colleagues and extended family. And in 1946 Hermann Haber died by suicide, reportedly because of his shame over his father’s wartime research.
How does one weigh the life of such a person?
In an address to Parliament in 1918, Winston Churchill remarked:
It is a very strange thing to reflect that but for the invention of Professor Haber the Germans could not have continued the War after their original stack of nitrates was exhausted. The invention of this single man has enabled them … not only to maintain an almost unlimited supply of explosives for all purposes, but to provide amply for the needs of agriculture in chemical manures. It is a remarkable fact, and shows on what obscure and accidental incidents the fortunes possibly of the whole world may turn in these days of scientific discovery.
If you go strictly by the numbers, perhaps Haber’s effects on recent history yielded a net positive. More than one hundred million war dead and great devastation throughout the last century might in theory be counterbalanced by the effects of air-derived fertilizer on human hunger. It has been estimated that as many as half of today’s seven billion people couldn’t even exist if it weren’t for artificially fixed nitrogen because there simply would not be enough usable nitrogen atoms available to build and maintain their bodies with. The outcomes of such equations depend on how the factors are weighted, and it could also be argued that more people can mean more pollution and strife in the world.
However you judge it, the Haber-Bosch process certainly liberated us from the constraints of the ancient bacterial cartel. Global commercial ammonia production currently exceeds one hundred million tons per year, and like other nitrogen fixers we also pass our nitrogenous wastes out into our surroundings in many forms. Thanks to the science and technology of a modern civilization that Haber helped to create, we can now explore those atomic connections in unprecedented and fascinating detail.
* * *
It is often said that “you are what you eat.” Some pioneering ecologists are now demonstrating this principle in the wilds of North America’s Pacific Northwest. Their primary subjects are not people, however, but salmon.
The annual migration of sockeye salmon upriver from the sea to spawn and die in remote lakes is a spectacular example of determination in the face of adversity. It has also been much more than that to residents of forested watersheds along that rugged coast for thousands of years. Salmon, quite literally, are part of the fabric of life there. And nitrogen, or more specifically a heavy stable isotope of it, is allowing scientists to tell new versions of that ancient story.
When you see a river seething with salmon, it is easy to lose track of individuals amid the chaos. Because of this problem, biologists use marker tags to track single fish in order to work out where they and their companions go, how long they live, and how many are caught relative to the population as a whole. But if you want to dig deeper into the story of salmon, normal labels won’t suffice. When a fish dies, its tag falls off and is lost. A new kind of marker now enables us to follow a salmon’s nitrogen atoms far beyond the fish to the occupants of surrounding forests as well as the forests themselves.
According to one biologist at the University of Victoria, British Columbia, who studies the atomic nature of salmon for a living, it all starts with the smell of rotting seaweed.
“You know that odor that hits you when you walk down to the beach,” Tom Reimchen said when I spoke to him by phone. “When seaweed decomposes it releases ammonia and other gases as bacteria break it down. This sort of thing also happens in the rest of the ocean whenever something decays.” For Reimchen and his colleagues, the isotopic composition of those molecules permits the tracking of salmon atoms through ecosystems.
“There are two main kinds of nitrogen atom out there,” he continued. “One of them is a little heavier than the other, and it escapes into the air a little less easily. This difference makes the heavier isotope build up in the oceans.”
Although you might think that atoms are as uniform as a school of fish seems to be, the existence of isotopes shows otherwise. Nitrogen-15 carries one more neutron than N-14, and being a little overweight, N-15 is a slightly more reluctant traveler than its lighter cousin. To people like Reimchen that reluctance is a good thing.
“Heavy N-15 vibrates less than N-14, so when it binds to other atoms the bonds tend to be more stable and harder to break. The ammonia rising out of rotten seaweed is enriched in N-14 because microbes have an easier time breaking it loose from proteins, which leaves more N-15 behind.” This also happens in marine food chains, with every step from prey to predator increasing the N-15 load of the organism. Single-celled algae consume the nitrogen compounds of dead seaweed, and the process of converting those food molecules into algal cells concentrates the N-15 relative to N-14. Tiny planktonic creatures eat the algae and concentrate the N-15 some more, small fish eat them, and so on. Sockeye salmon exist near the top of their food chain, and they feed only in the N-15 enriched oceans, so they carry distinctively large fractions of N-15 in their bodies. Scientists routinely find N-15 concentrations more than ten times higher than atmospheric or terrestrial values within the tissues of sockeyes.
“A fish that feeds in the ocean contains more N-15 than a freshwater fish does,” Reimchen explained, “and the same is true of anything else that eats food from the sea rather than from the land or a typical river or lake.” As a result, high concentrations of N-15 in body tissues can reveal subtle atomic ties to the ocean in a previously undocumented ecosystem, the “salmon forest.”
Shape-shifting has long been a mainstay of mythology throughout the world. In such traditions witches become black cats, hairy men become werewolves, and vampires become bats. We may chuckle at those tales, but scientists such as Reimchen don’t blink an eye. In the realm of atoms, fish become bears, wolves, birds, and human beings, and deep oceans can become lush forests.
The Queen Charlotte Islands (more formally known as Haida Gwaii) lie just off the coast of British Columbia, and they attract migratory salmon into their rivers year after year. In an article published in Ecoforestry, Reimchen reported that a single Queen Charlotte black bear can capture as many as seven hundred salmon during the six-week spawning period in late summer and early autumn. Normally a bear carries its prize off into the woods to dine in private, and it tends to eat mostly the choicest bits, leaving the rest of the carcass on the ground while it returns to the hunt. In the realm of atoms, food rarely goes to waste, and the leftovers of bears’ feasts soon become dinner for eagles, ravens, and gulls. Days later any uneaten soft parts and animal droppings are bonanzas for beetles, flies, and bacteria that turn them into nitrogenous wastes. The decaying refuse in the riverside forest produces, as Reimchen put it, “a highly odiferous riparian zone.”
As with black bears, a grizzly’s nitrogen intake can also come almost exclusively from fish protein during the annual salmon migration in British Columbia. After a lunge, a clamping of jaws, and a shaking of heavy wet fur, a successful hunter drags its prey into the bushes. Tearing open the skin to expose orange-pink meat, the bear gulps a mouthful of fish protein into its stomach, where enzymes chop it into smaller, nitrogen-bearing amino acids. Amino acids from the meal ride the predator’s bloodstream to the liver, where most of them are processed into the raw materials for muscles, sinews, and other body parts. Anything above the daily requirement for growth, tissue replacement, and such things is broken down into ammonia, which is then converted into urea for disposal. Nine out of ten of those discarded nitrogen compounds are filtered out by the kidneys and excreted in urine, and most of the rest exit through solid waste. Similar things would also happen inside you if you should likewise take a bite of savory salmon and make those fish atoms your own.
If the bear is a lactating mother, then much of the unused nitrogen goes into milk protein instead, but the body normally uses most of the amino acids for producing fur and other items of bear anatomy. In one study 80 percent of the nitrogen in bear fur could be traced back to the ocean. This atomic signal was present in newly emerging hair shafts not only during the fall salmon migration, but also in spring and early summer, when the bears were eating more vegetation than fish. Oceanic nitrogen, it turns out, suffuses the entire forest, not just the bears.
Plants use nitrogen atoms in the light-trapping chlorophyll that turns them green, as well as in enzymes, cell walls, and fragrances. Near salmon streams much of the nitrogen that plants absorb from the soil comes from bear-processed fish proteins, and Reimchen and his colleagues have measured high concentrations of N-15 in local huckleberries, wild azaleas, and the aptly named salmonberries, which reveal in atomic script the isotopic signature of the sea. Similar studies have found the same pattern in watersheds all along the northern Pacific Coast, in which the N-15 content of foliage reflects the proximity of plants to a stream, the density of fish, and the abundance of bears.
Unfortunately, less desirable gifts from the sea also travel through these food webs with the nitrogen atoms. According to a study by the environmental scientist Jennie Christensen and her colleagues, airborne pollutants that fall into the Pacific Ocean can eventually contaminate salmon, as well. They concluded that migratory salmon deliver up to 70 percent of the organochlorine pesticides and 90 percent of the toxic PCBs found in fish-eating grizzlies, “thereby inextricably linking these terrestrial predators to contaminants from the North Pacific Ocean.”
Meanwhile, back at the carcass, other bits of salmon nitrogen are trotting away in the stomachs of wolves, weasels, and foxes, or flying up into the trees with the eagles and ravens. Some make detours through carrion-feeding insects to flycatchers and warblers before falling back to the ground in bird droppings. Even the nitrogen atoms in the leftover fish skeletons slowly escape as bacteria release ammonia and nitrate into the soil, where fungal fibers distribute them among the roots of the forest.
The trunks of the spruce and hemlocks lining Reimchen’s study streams contain more N-15 than those of trees growing farther upslope or upstream of waterfalls that block salmon migrations. As much as 40 percent of the nitrogen in riverside vegetation at the British Columbia sites may come from the bodies of fish by way of bears, but at one coastal stream where the highest salmon densities were found, Reimchen estimated that some of the heaviest migration years see more than three-quarters of the annual nitrogen budgets of the local spruce trees supplied by fish.
Studies show that salmon nitrogen can make plants grow faster or more profusely, and some investigators have used tree rings as recorders of salmon abundance through time. Most fisheries studies cover only a few years or decades, but spruce, fir, and hemlock trees can live for centuries, laying down annual bands of new wood in the outer layers of their trunks. By measuring the abundance of N-15 ring by ring in streamside trees, researchers from the University of Washington found that salmon populations fluctuated a great deal during the last 350 years, even before modern human impacts began. But the longest records of salmon history come from lakes in which survivors of the predator gauntlet end their epic journeys.
As sockeyes work their way up into the headwaters of their home streams, the males lose their silvery sheen and develop red flanks, green heads, and wickedly curved jaws in preparation for courtship and turf battles. When the fish reach a lake, they begin to spawn. Females lay eggs in the shallows, and males fight for a chance to drop their milky gametes on the nests before dying. Soon the lake floor is littered with battered carcasses, which decompose and release nitrogen atoms into an aquatic version of the salmon forests.
Algae consume the nitrogen compounds and then become food for microscopic zooplankton, producing bite-size meals for newly hatched salmon. The youngsters spend several years fattening up on the recycled atoms of their forebears before making their own round trips to and from the Pacific. And all the while, thin films of detritus and dead plankton accumulate in the soft mud beneath the lake like annual bands in a tree trunk. These layered sediment archives can represent thousands of years of salmon history.
In a paper published in Nature in 2002, ecologist Bruce Finney and his colleagues described what they found in sediment cores from Karluk Lake, a wild sockeye nursery on Kodiak Island, Alaska. During the last two millennia, sedimentary N-15 values rose and fell along with the glassy shells of diatom algae that tend to grow most prolifically in years when salmon migrations are largest. But more surprising than the length and detail of the records was what they revealed about today’s populations.
Long-term declines such as those that are now under way are not unique to our times, and an even larger crash that occurred two thousand years ago lasted for several centuries. We don’t yet know what caused it, although Finney’s team suspected that it had something to do with fluctuations in climates, Pacific currents, and sea surface temperatures.
Many other questions also emerge from these discoveries. How abundant are salmon “supposed” to be? Were old reports of rivers so full that you could cross on the backs of the fish just “big fish” tales? When commercial fishing first started in the region, salmon were more abundant than they had been for thousands of years. Did that coincidence leave us with a false impression of what to expect in the future?
Such investigations show that we still have a lot to learn about salmon and their place in the nitrogen economy of the Pacific Northwest. And not surprisingly, similar studies are beginning to uncover a great deal about our own nitrogenous connections to the world as well.
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Through most of human history, the atomic exchanges between our ancestors and their surroundings were much like those in a salmon forest, and they can still be traced through human bodies with isotopic markers just as they are in Reimchen’s research.
Archaeologists are using N-15 analyses of collagen, the most abundant protein in bone, to reconstruct the diets of people who died long ago. In one such study, the archaeologist Erle Nelson and his colleagues exhumed the skeletons of early Norse settlers in southern Greenland in order to help determine why their settlements failed during the fifteenth century. Some experts have proposed that a natural cooling event, the so-called Little Ice Age, made conditions too harsh. Others have surmised that a stubborn refusal to adopt the hunting and fishing habits of the local Inuit led to death by soil depletion and crop failures. In 2012 Nelson’s team probed the atomic contents of Greenlandic graveyards in order to test these competing ideas.
Their isotopic analyses showed that the Greenlandic Norse ate quite a bit of fish and seal meat, thereby weakening claims that climate or crop failures did them in. As with the bears of the Pacific Northwest, people who eat marine creatures carry more N-15 within them than do people whose nitrogen atoms come from crops and domestic livestock. In the Greenland study all the skeletons contained more N-15 than they would have if the settlers had lived, as it were, on bread alone.
All the skeletons, that is, except one. According to local documents, this individual was born in Scandinavia and had joined the colonists only a short time before he died. As a result his bones still carried many of the agricultural nitrogen atoms from meals that he consumed during his earlier years back in Norway. The relatively low N-15 contents of his remains confirmed that he had not yet replaced all of the nitrogen in his body with local food atoms at the time of his death.
According to the authors, the nitrogen isotopes of those early settlers showed that the demise of the Greenland colonies was probably not so much a sudden collapse as a slow drain-off. In a press release from the University of Copenhagen, the anthropologist Niels Lynnerup explained: “Nothing suggests that the Norse disappeared as a result of a natural disaster. If anything, they might have become bored with eating seals out on the edge of the world.” Longing for the richer cultural and social environments of Scandinavia may simply have lured so many young people away that the remaining villagers finally gave up and left, too.
Digesting the proteins of another organism tends to leave more of the heavy N-15 atoms behind in one’s tissues, and this typically makes the N-15 concentration of a predator’s body higher than that of its prey. The same principle applies to us. According to a paper published in Organic Geochemistry in 1997, a wonderful manifestation of this kind of atomic linkage can be found in the nutritional connections between nursing mothers and their infants.
When you were still a fetus, all your atoms came from your mother’s body. No eating, drinking, or breathing for you yet, only the cyclic pumping of fluids between placenta and umbilical cord. You were essentially a piece of your mom then, and although you breathed your own oxygen after you were born, you would still have gotten your other atoms through mother’s milk during infancy unless you were raised on bottled formula.
But although your first atoms came from your mother, the body that you built from those atoms was slightly enriched with N-15 in much the same manner that a bear’s proteins are enriched relative to the salmon it eats. In other words, you were not only a part of your mother: You were also feeding on her. Call this relationship predation, parasitism, or even cannibalism if you like; it was all the same to your nitrogen isotopes, which recorded what you were doing and wrote it into the molecular archives of your proteins.
When investigators gently trimmed the fingernails of breast-fed newborns and compared them to fingernail clippings from their mothers, they found the telltale isotopic signs of that maternal sacrifice. A captured salmon unwillingly feeds its N-15 to a hungry bear, but nursing mothers freely offer their own atoms to their babies. In keeping with the age-old rule of food chains, the proteins of the infants in the study were a little richer in N-15 than those of their moms, and they remained so until the children were weaned and their isotopic contents shifted into closer alignment with the atoms of the larger world that would henceforth sustain them. The first isotopic signs of independence appeared in clippings after two to three months, the time it takes for a newborn’s fingernail to grow from cuticle to fingertip. Similar isotopic studies show that the course of a pregnancy can also be traced along the lengths of a mother’s hair filaments as nitrogen atoms drain from her body into her child.
In many ways the intimate connections between mother and child are also mirrored in our atomic ties to the Earth. Even after we pull away from our early umbilical linkage, we still live embedded in the same global reservoir of recycled atoms. But today those atomic connections join us not only to living things but also to machines.
At the time of this writing, our vehicles, farms, and industries produce about half of the biologically useful nitrogen on Earth. One study published in Science in 2010 reported that the Haber-Bosch process alone now matches the entire nitrogen-fixation output of the oceans and exceeds the microbial fixation on land. We live in a world that is fundamentally different from the ones that our distant ancestors knew. No, most of us are not in Kansas anymore.
The people in the black-and-white world of The Wizard of Oz lived off whatever their farm could provide, and the nitrogen in their bodies came from their livestock, crops, and the occasional lightning bolt. From soil bacteria to plants and farm residents and back again, nitrogen atoms returned over and over to the same soil as horse teams plowed it and the hired help spread hog droppings and other wastes on it. Not a perfectly idyllic existence by any means in those hard times, even without a twister about to strike, but one from which we still have important lessons to relearn in greater depth. Those people were part of an elemental cycle that is now largely broken in present-day Kansas as well as in much of the rest of the world.
If you are like most Americans, you obtain nearly all of your food from supermarkets that are supplied by industrial-scale farms. It would therefore be safe to assume that most of the nitrogen atoms in your own body were artificially fixed a long distance from where you live. Most of the energy for that fixation came from massive amounts of nonrenewable fossil fuel, as did the energy for the transportation of the food to your local stores and then to your table. Many of those costly nitrogen compounds were never absorbed by crops in the fields upon which they were sprayed, and others escaped from livestock in wastes that ended up in storage lagoons or sewage treatment facilities where denitrifying bacteria returned them to the air as inert nitrogen gas. Such processes open enormous gaps in the human-centered portions of the global nitrogen cycle that must be filled by yet more energy-intensive fixation and transport.
One could argue that this is how it must be, that this is now the only way to keep so many people alive and free from the agony of hunger. But finding ways to more effectively close the loop on our nitrogen cycle is both feasible and beneficial in the long run. Returning animal wastes to the soil, replenishing fields with nitrogen-fixing crops, and more carefully adjusting the timing and amounts of fertilizers that we spread on our fields are among the many relatively simple suggestions now in circulation.
For hundreds of thousands of years our ancestors absorbed, used, and released these same atoms back into the same finite pool, trusting the world to provide them with what they needed and to dilute or dispose of their wastes. The Haber-Bosch process has propelled us into this present century as the dominant source of living nitrogen on Earth, and as our ability to draw nourishment from the atmosphere increases, so too does our influence on the global nitrogen cycle.
The ecologist David Schindler described our situation succinctly in an interview for the University of Washington. “The world, for nitrogen,” he said, “is a much smaller place than we’d assumed.” The nitrogen atoms that we use remain on the planet with us when we discard them, and they now do so in ever-growing quantities along with our other wastes. In addition to feeding our crops, our nitrogenous emissions also trigger unwanted algae blooms, acidify rain and snow, and help to toxify urban smog. They even contribute to climate change, because the nitrous oxide that is released from overfertilized fields, lawns, golf courses, and fossil fuel combustion is a greenhouse gas. As the ecologist Alex Wolfe noted in an online interview, “The global change debate is dominated by discussions of carbon emissions … [but] the global nitrogen cycle has been far more perturbed by humanity than that of carbon.”
How we balance our need for fixed nitrogen against its effects on air and water quality, as well as our evolving ability to feed or to kill one another, will be one of the great unfolding stories of our journey forward as a rapidly maturing, sentient species. And however that story develops in the future, the ambiguous legacy of Fritz Haber will surely continue to play a role in it, too.